Functionalized Inorganic Fluorides: Synthesis, Characterization and Properties of Nanostructured Solids

Functionalized Inorganic Fluorides Functionalized Inorganic Fluorides Synthesis, Characterization & Properties of Nan...

Author: Alain Tressaud

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Functionalized Inorganic Fluorides

Functionalized Inorganic Fluorides Synthesis, Characterization & Properties of Nanostructured Solids

Edited by ALAIN TRESSAUD Research Director CNRS (Emeritus), ICMCB–CNRS, Bordeaux University, France

A John Wiley and Sons, Ltd., Publication

This edition first published 2010 Ó 2010 John Wiley & Sons, Ltd Registered office John Wiley & Sons Ltd, The Atrium, Southern Gate, Chichester, West Sussex, PO19 8SQ, United Kingdom For details of our global editorial offices, for customer services and for information about how to apply for permission to reuse the copyright material in this book please see our website at www.wiley.com. The right of the author to be identified as the author of this work has been asserted in accordance with the Copyright, Designs and Patents Act 1988. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, except as permitted by the UK Copyright, Designs and Patents Act 1988, without the prior permission of the publisher. Wiley also publishes its books in a variety of electronic formats. Some content that appears in print may not be available in electronic books. Designations used by companies to distinguish their products are often claimed as trademarks. All brand names and product names used in this book are trade names, service marks, trademarks or registered trademarks of their respective owners. The publisher is not associated with any product or vendor mentioned in this book. This publication is designed to provide accurate and authoritative information in regard to the subject matter covered. It is sold on the understanding that the publisher is not engaged in rendering professional services. If professional advice or other expert assistance is required, the services of a competent professional should be sought. The publisher and the author make no representations or warranties with respect to the accuracy or completeness of the contents of this work and specifically disclaim all warranties, including without limitation any implied warranties of fitness for a particular purpose. This work is sold with the understanding that the publisher is not engaged in rendering professional services. The advice and strategies contained herein may not be suitable for every situation. In view of ongoing research, equipment modifications, changes in governmental regulations, and the constant flow of information relating to the use of experimental reagents, equipment, and devices, the reader is urged to review and evaluate the information provided in the package insert or instructions for each chemical, piece of equipment, reagent, or device for, among other things, any changes in the instructions or indication of usage and for added warnings and precautions. The fact that an organization or Website is referred to in this work as a citation and/or a potential source of further information does not mean that the author or the publisher endorses the information the organization or Website may provide or recommendations it may make. Further, readers should be aware that Internet Websites listed in this work may have changed or disappeared between when this work was written and when it is read. No warranty may be created or extended by any promotional statements for this work. Neither the publisher nor the author shall be liable for any damages arising herefrom. Library of Congress Cataloging-in-Publication Data Functionalized inorganic fluorides: synthesis, characterization & properties of nanostructured solids / edited by Alain Tressaud. p. cm. Includes bibliographical references and index. ISBN 978-0-470-74050-7 (pbk.) 1. Fluorides. I. Tressaud, Alain. QD181.F1F77 2010 5460 .731—dc22 2009052139 A catalogue record for this book is available from the British Library. ISBN: 978-0-470-74050-7 (Cloth) Set in 10/12pt Times by Integra Software Services Pvt. Ltd., Pondicherry, India Printed and bound in Great Britain by CPI Antony Rowe, Chippenham, Wiltshire. Cover images from left to right: Projection along [001] of the ITQ-33 zeolite structure showing the 18-MRs windows (Chapter 16); Schematic morphology of oxyfluoride glass-ceramics formed by spinodal decomposition (Chapter 9); Crystal structure of La2CuO3.6F0.8 [The Cu cations are situated in octahedra; the La cations are shown as large spheres; the F anions are shown as small spheres] (Chapter 13)

Contents

Preface List of Contributors 1

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties Erhard Kemnitz, Gudrun Scholz and Stephan Ru¨diger

1

1.1

1

Introduction 1.1.1 Sol-Gel Syntheses of Oxides – An Intensively Studied and Widely Used Process 1.1.2 Sol-Gel Syntheses of Metal Fluorides – Overview of Methods 1.2 Fluorolytic Sol-Gel Synthesis 1.2.1 Mechanism and Properties 1.2.2 Insight into Mechanism by Analytical Methods 1.2.3 Exploring Properties 1.2.4 Possible Fields of Application References 2

xvii xxi

1 2 4 5 8 27 29 35

Microwave-Assisted Route Towards Fluorinated Nanomaterials Damien Dambournet, Alain Demourgues and Alain Tressaud

39

2.1 2.2

39 40 40 40 41 41 42 42

2.3

Introduction Introduction to Microwave Synthesis 2.2.1 A Brief History 2.2.2 Mechanisms to Generate Heat 2.2.3 Advantages of Microwave Synthesis 2.2.4 Examples of Microwave Experiments Preparation of Nanosized Metal Fluorides 2.3.1 Aluminium-based Fluoride Materials 2.3.2 Microwave-assisted Synthesis of Transition Metal Oxy-Hydroxy-Fluorides

61

vi

3

Contents

2.4 Concluding Remarks Acknowledgements References

64 64 65

High Surface Area Metal Fluorides as Catalysts Erhard Kemnitz and Stephan Ru¨diger

69

3.1 3.2 3.3 3.4 3.5 3.6

69 71 74 78 84 88 90 94 95 97

Introduction High Surface Area Aluminium Fluoride as Catalyst Host-Guest Metal Fluoride Systems Hydroxy(oxo)fluorides as Bi-acidic Catalysts Oxidation Catalysis Metal Fluoride Supported Noble Metal Catalysts 3.6.1 Hydrodechlorination of Monochlorodifluoromethane 3.6.2 Hydrodechlorination of Dichloroacetic Acid (DCA) 3.6.3 Suzuki Coupling References

4

Investigation of Surface Acidity using a Range of Probe Molecules Alexandre Vimont, Marco Daturi and John M. Winfield

101

4.1

101 102

Introduction 4.1.1 Setting the Scene: Metal Fluorides versus Metal Oxides 4.1.2 Some Examples of the Application of FTIR Spectroscopy to the Study of Surface Acidity in Metal Oxides 4.1.3 A Preview 4.2 Characterization of Acidity on a Surface: Contrasts with Molecular Fluorides 4.2.1 Molecular Brønsted and Molecular Lewis Acids 4.2.2 A Possible Benchmark for Solid Metal Fluoride, Lewis Acids: Aluminium Chlorofluoride 4.3 Experimental Methodology 4.3.1 FTIR Spectroscopy 4.3.2 Characteristic Reactions and the Detection of Adsorbed Species by a Radiotracer Method 4.4 Experimental Studies of Surface Acidity 4.4.1 Using FTIR Spectroscopy 4.4.2 Using HCl as a Probe with Detection via [36Cl]-Labelling 4.4.3 Metal Fluoride Surfaces that Contain Surface Hydroxyl Groups: Aluminium Hydroxy Fluorides with the Hexagonal Tungsten Bronze Structure 4.4.4 Possible Geometries for HCl Adsorbed at Metal Fluoride Surfaces: Relation to Oxide and Oxyfluoride Surfaces 4.5 Conclusions References

103 107 108 108 109 110 110 112 117 118 123

129 135 136 137

Contents

5

Probing Short and Medium Range Order in Al-based Fluorides using High Resolution Solid State Nuclear Magnetic Resonance and Parameter Modelling Christophe Legein, Monique Body, Jean-Yves Buzare´, Charlotte Martineau and Gilles Silly 5.1 5.2

Introduction High Resolution NMR Techniques 5.2.1 Fast MAS and High Magnetic Field 27 Al NMR 5.2.2 5.2.3 High Resolution Correlation NMR Techniques 5.3 Application to Functionalized Al-Based Fluorides with Catalytic Properties 5.3.1 Crystalline Aluminium Fluoride Phases 19 F Isotropic Chemical Shift Scale in Octahedral Aluminium 5.3.2 Environments with Oxygen and Fluorine in the First Coordination Sphere 5.3.3 Fluorinated Aluminas and Zeolites, HS AlF3 5.3.4 Aluminium Chlorofluoride and Bromofluoride 5.3.5 Pentahedral and Tetrahedral Aluminium Fluoride Species 5.3.6 Nanostructured Aluminium Hydroxyfluorides and Aluminium Fluoride Hydrate with Cationic Vacancies 5.3.7 iso Scale for 27Al and 19F in Octahedral Aluminium Environments with Hydroxyl and Fluorine in the First Coordination Sphere 5.4 Alkali and Alkaline-earth Fluoroaluminates: Model Compounds for Modelling of NMR Parameters 19 F NMR Line Assignments 5.4.1 27 Al Site assignments, Structural and Electronic 5.4.2 Characterizations 5.5 Conclusion References

6

vii

141

141 142 142 145 148 153 153

153 157 158 158

159

160 160 161 164 167 168

Predictive Modelling of Aluminium Fluoride Surfaces Christine L. Bailey, Sanghamitra Mukhopadhyay, Adrian Wander, Barry Searle and Nicholas Harrison

175

6.1 6.2

175 176 176 177 178 179 180 180 180

6.3

Introduction Methodology 6.2.1 Density Functional Theory 6.2.2 Surface Free Energies 6.2.3 Molecular Adsorption 6.2.4 Kinetic Monte Carlo Simulations Geometric Structure of and -AlF3 6.3.1 Bulk Phases 6.3.2 Surfaces

viii

Contents

6.4 6.5

Characterization of AlF3 Surfaces Surface Composition under Reaction Conditions 6.5.1 The -AlF3–x (01–12) Termination 6.5.2 The -AlF3 (0001) Termination 6.6 Characterization of Hydroxylated Surfaces 6.7 Surface Catalysis 6.7.1 Molecular Adsorption 6.7.2 Reaction Mechanisms and Barriers 6.7.3 Analysing the Kinetics of the Reaction 6.8 Conclusions Acknowledgements References

7

Inorganic Fluoride Materials from Solvay Fluor and their Industrial Applications Placido Garcia Juan, Hans-Walter Swidersky, Thomas Schwarze and Johannes Eicher 7.1 7.2

7.3

7.4 7.5

7.6 7.7 7.8 7.9 7.10 7.11 7.12 7.13 7.14 7.15 7.16 7.17

Introduction Hydrogen Fluoride 7.2.1 Anhydrous Hydrogen Fluoride, AHF 7.2.2 Hydrofluoric Acid Elemental Fluorine, F2 7.3.1 Fluorination of Plastic Fuel Tanks 7.3.2 Finishing of Plastic Surfaces 7.3.3 F2 Mixtures as CVD-chamber Cleaning Gas Iodine Pentafluoride, IF5 Sulfur Hexafluoride, SF6 7.5.1 SF6 as Insulating Gas for Electrical Equipment 7.5.2 SF6 Applications in Metallurgy Ammonium Bifluoride, NH4HF2 Potassium Fluorometalates, KZnF3 and K2SiF6 Cryolite and Related Hexafluoroaluminates, Na3AlF6, Li3AlF6, K3AlF6 Potassium Fluoroborate, KBF4 Fluoboric Acid, HBF4 Barium Fluoride, BaF2 Synthetic Calcium Fluoride, CaF2 Sodium Fluoride, NaF Sodium Bifluoride, NaHF2 Potassium Bifluoride, KHF2 Potassium Fluoroaluminate, KAlF4 Fluoroaluminate Fluxes in Aluminium Brazing 7.17.1 Flux Composition 7.17.2 Flux and HF 7.17.3 Flux Particle Size

185 188 189 192 193 196 197 198 200 201 203 203

205

205 205 206 206 207 207 207 208 208 209 209 209 210 210 211 212 212 213 213 213 213 214 214 214 214 216 217

Contents

8

7.17.4 Flux Melting Range 7.17.5 Current Status of Aluminium Brazing Technology 7.17.6 Cleaning and Flux Application 7.17.7 Wet Flux Application 7.17.8 Dry/Electrostatic Flux Application 7.17.9 Post Braze Flux Residue 7.17.10 Filler Metal Alloys 7.17.11 Flux Precoated Brazing Sheet/Components 7.17.12 Clad-less Brazing 7.17.13 Furnace Conditions 7.18 Summary References

219 220 221 221 222 222 222 223 223 224 224 225

New Nanostructured Fluorocompounds as UV Absorbers Alain Demourgues, Laetitia Sronek and Nicolas Penin

229

8.1 8.2

229 231 231 232

Introduction Synthesis of Tetravalent Ce and Ti-based Oxyfluorides 8.2.1 Preparation of Ce-Ca-based Oxyfluorides 8.2.2 Preparation of Ti-based Oxyfluorides 8.3 Chemical Compositions and Structural Features of Ce and Ti-based Oxyfluorides 8.3.1 Elemental Analysis 8.3.2 Magnetic Measurements 8.3.3 About the Chemical Composition of Ce1xCaxO2x and Ce1xCaxO2xy/2Fy Series 8.3.4 About the Structure and Local Environment of Fluorine in Ce1xCaxO2xy/2Fy Series 8.3.5 Composition and Structure of Ti-based Hydroxyfluoride 8.4 UV Shielding Properties of Divided Oxyfluorides 8.4.1 The Ce-Ca-based Oxyfluorides Series and UV-shielding Properties 8.4.2 Ti Hydroxyfluoride and UV-shielding Properties 8.5 Conclusion Acknowledgement References 9

ix

233 233 233 234 237 252 263 264 266 267 269 269

Oxyfluoride Transparent Glass Ceramics Michel Mortier and Ge´raldine Dantelle

273

9.1 9.2

273 274 275 277 279 281 281

9.3

Introduction Synthesis 9.2.1 Synthesis by Glass Devitrification 9.2.2 Transparency Different Systems 9.3.1 Glass-Ceramics with CaF2 as their Crystalline Phase 9.3.2 Glass-Ceramics with -PbF2 as their Crystalline Phase

x

Contents

9.3.3 Glass-Ceramics with CdF2/PbF2 as their Crystalline Phase 9.3.4 Glass-Ceramics with LaF3 as their Crystalline Phase 9.4 Thermal Characterization 9.4.1 Kinetics of Phase-change/Devitrification 9.4.2 Thakur’s Method 9.5 Morphology of the Separated Phases 9.6 Optical Properties of Glass-Ceramics 9.6.1 Influence of the Devitrification on the Spectroscopic Properties of Ln3þ 9.6.2 Effect of High Local Ln3þ Concentration in Crystallites 9.6.3 Comparison of the Optical Properties of Glass-Ceramics and Single-Crystals 9.6.4 Multi-doped Glass-Ceramics 9.7 Conclusion References 10

11

Sol-Gel Route to Inorganic Fluoride Nanomaterials with Optical Properties Shinobu Fujihara

281 282 282 288 288 289 293 293 295 297 299 301 302

307

10.1 10.2

Introduction Principles of a Sol-Gel Method 10.2.1 Metal Oxide Materials 10.2.2 Metal Fluoride Materials 10.3 Fluorinating Reagents and Method of Fluorination 10.4 Control of Shapes and Microstructures 10.5 Optical Properties 10.5.1 Low Refractive Index and Anti-Reflection Effect 10.5.2 Luminescence 10.6 Concluding Remarks References

307 308 308 308 309 313 317 317 322 326 326

Fluoride Glasses and Planar Optical Waveguides Brigitte Boulard

331

11.1 11.2

331 332 333 334 336 338 340 341 342 344 344

Introduction Rare Earth in Fluoride Glasses 11.2.1 Fundamentals 11.2.2 Applications: Laser and Optical Amplifiers 11.3 Fabrication of Waveguides: A Review 11.4 Performance of Active Waveguides 11.4.1 Optical Amplifier 11.4.2 Lasers 11.5 Fluoride Transparent Glass Ceramics: An Emerging Material 11.6 Conclusion References

Contents

12

13

xi

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates Karim Adil, Amandine Cadiau, Annie He´mon-Ribaud, Marc Leblanc and Vincent Maisonneuve

347

12.1 12.2 12.3

Introduction Synthesis Extended Finite Polyanions (0D) 12.3.1 Isolated AlF4 Tetrahedra 12.3.2 Isolated AlF6 Octahedra 12.3.3 Al2F11 Dimers 12.3.4 Al3F16 Trimers 12.3.5 Al2F10 Dimers 12.3.6 Al4F20 Tetramers 12.3.7 Al4F18 Tetramers 12.3.8 Al5F26 Pentamers 12.3.9 Al7F30 Heptamers 12.3.10 Al8F35 Octamers 12.3.11 Mixed Polyanions 12.4 1D Networks 12.4.1 AlF5 Chains 12.4.2 Al2F9 Chains 12.4.3 Al7F29 Chains 12.4.4 AlF4 Chains 12.4.5 Mixed Polyanions and/or Chains 12.5 2D Networks 12.5.1 Al3F14 Layers 12.5.2 AlF4 Layers 12.5.3 Al2F7 Layers 12.5.4 Al5F17 Layers 12.5.5 Al3F10 Layers 12.6 3D Networks 12.6.1 Al7F33 Network 12.6.2 Al2F9 Network 12.6.3 AlF3 Network 12.7 Evolution of the Condensation of Inorganic Polyanions 12.7.1 Influence of Amine and Aluminum Concentrations 12.7.2 Temperature Acknowledgements Supplementary Materials References

347 348 350 350 350 353 353 353 354 354 355 355 356 356 358 358 359 360 360 361 365 365 365 366 366 368 368 368 368 369 372 372 374 376 376 376

Synthesis, Structure and Superconducting/Magnetic Properties of Cu- and Mn-based Oxyfluorides Evgeny V. Antipov and Artem M. Abakumov

383

13.1 13.2

383 384

Introduction Chemical Aspects of Fluorination of Complex Oxides

xii

Contents

13.3

Structural Aspects of Fluorination of Complex Cuprates and Superconducting Properties 13.3.1 Electron Doped Superconductors: Heterovalent Replacement 1O2 ! 1F 13.3.2 Hole Doped Superconductors: Fluorine Insertion into Vacant Anion Sites 13.3.3 Structural Rearrangements in Fluorinated Cuprates 13.3.4 Fluorination of Nonsuperconducting Cuprates 13.4 Fluorination of Manganites 13.5 Conclusions References 14

Doping Influence on the Defect Structure and Ionic Conductivity of Fluorine-containing Phases Elena I. Ardashnikova, Vladimir A. Prituzhalov and Ilya B. Kutsenok 14.1 14.2

14.3

14.4 14.5

14.6

14.7

14.8

Introduction Influence of Oxygen Ions on Fluoride Properties 14.2.1 Pyrohydrolysis 14.2.2 Heterovalent Oxygen Substitution for Fluoride Ions 14.2.3 Ionic Conductivity of Oxyfluoride Cation Doping of Fluorides 14.3.1 Isovalent Replacement in the Cation Sublattice 14.3.2 Heterovalent Replacement in the Cation Sublattice Active Lone Electron Pair of Cations and Ionic Conductivity Peculiarities of the Defect Structure of Nonstoichiometric Fluorite-like Phases 14.5.1 Fluorite Structure 14.5.2 Defect Clusters 14.5.3 Ordered Fluorite-like Phases 14.5.4 Phase Diagrams Ionic Transfer in Fluorite-like Phases 14.6.1 Defect Region Model 14.6.2 Nonstoichiometric Fluorites as Examples of Nanostructured Materials Peculiarities of the Defect Structure of Nonstoichiometric Tysonite-like Phases 14.7.1 Tysonite Structure, Tysonite Modifications and Anion Defects 14.7.2 Ordered Tysonite-like Phases Ionic Transfer in Tysonite-like Phases 14.8.1 Fluoride Ions’ Migration Paths in the LaF3 Structure 14.8.2 Temperature Dependences of Ionic Conductivity and Anion Defect Positions 14.8.3 Concentration Dependences of Ionic Conductivity in Tysonite-like Solid Solutions

388 389 390 398 408 411 415 416

423 423 427 427 428 429 431 431 432 432 435 435 435 439 441 441 443 447 449 449 454 454 455 457 459

Contents

15

14.9 Conclusions References

462 462

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups Yoshiaki Matsuo

469

15.1 15.2

469

Introduction Preparation and Properties of Intercalation Compounds Containing Perfluoroalkyl Groups 15.2.1 Preparation 15.2.2 Exfoliation and Film Preparation 15.2.3 Introduction of Photofunctional Molecules 15.3 Photophysical and Photochemical Properties of Dyes in Intercalation Compounds Containing Perfluoroalkyl Groups 15.3.1 Microenvironment Estimated by using Probe Molecules Showing Photophysical Responses 15.3.2 Photophysical Properties 15.3.3 Photochemical Properties 15.4 Conclusion and Future Perspectives References 16

17

xiii

The Fluoride Route: A Good Opportunity for the Preparation of 2D and 3D Inorganic Microporous Frameworks Jean-Louis Paillaud, Philippe Caullet, Jocelyne Brendle´, Ange´lique Simon-Masseron and Joe¨l Patarin

471 471 475 476 478 478 480 482 484 484

489

16.1 16.2 16.3 16.4 16.5

Introduction Silica-based Microporous Materials Germanium-based Microporous Materials Phosphate-based Microporous Materials Synthetic Clays 16.5.1 Semi-Synthesis 16.5.2 Solid State Synthesis 16.5.3 Hydrothermal Synthesis 16.6 Conclusion References

489 490 499 504 506 507 508 509 510 511

Access to Highly Fluorinated Silica by Direct F2 Fluorination Alain Demourgues, Emilie Lataste, Etienne Durand and Alain Tressaud

519

17.1 17.2 17.3

519 520

17.4

Introduction Mesoporous Silica and Fluorination Procedures About the Chemical Composition and Morphology of Highly Fluorinated Silica FTIR Analysis 17.4.1 About the Content and Nature of OH/Water Groups in Highly Fluorinated Silica

521 523 523

xiv

Contents

17.4.2 17.4.3

FTIR Bands Related to Si-F Bonds Correlation between Silanol Groups on Mesoporous Silica and Grafted Fluorine on Highly Fluorinated Silica 17.5 Thermal Stability and Water Affinity of Highly Fluorinated Silica 17.6 Nuclear Magnetic Resonance (NMR) Investigations 17.6.1 Local Environments in Highly Fluorinated Silica through NMR Experiments 17.6.2 Effect of Fluorination on the Nuclei Environments 17.7 Conclusions on the F2-gas Fluorination Mechanism of Mesoporous Silica Acknowledgements References 18

Preparation and Properties of Rare-earth-Containing Oxide Fluoride Glasses Susumu Yonezawa, Jae-ho Kim and Masayuki Takashima 18.1 18.2

Introduction Preparation and Basic Characteristics of Oxide Fluoride Glasses Containing LnF3 18.2.1 Preparation of Oxide Fluoride Glasses Containing LnF3 18.2.2 Density and Refractive Index 18.2.3 Glass Transition Temperature 18.3 Optical and Magnetic Properties of LnF3-BaF2-AlF3-GeO2 (SiO2) Glasses 18.3.1 Optical Properties of HoF3-BaF2-AlF3-GeO2 Glasses 18.3.2 Optical Properties of CeF3-BaF2-AlF3-SiO2 Glasses 18.3.3 Optical Properties of the Glasses Co-doped with TbF3 and SmF3 18.3.4 Magnetic Property of TbF3 Containing Oxide Fluoride Glasses 18.4 Conclusion References

19

Switchable Hydrophobic-hydrophilic Fluorinated Layer for Offset Processing Alain Tressaud, Christine Labruge`re, Etienne Durand 19.1 19.2 19.3

19.4

Introduction The Principles of the Lithographic Printing Process Experimental Part 19.3.1 Fluorination by Cold rf Plasmas 19.3.2 Wettability Measurements 19.3.3 Surface Analyses Various Types of Surface Modifications using Fluorinated rf Plasmas 19.4.1 Reactive Etching of Porous Alumina using CF4-Plasma Treatment

526 527 530 533 534 534 540 541 542

545 545 546 546 552 553 555 555 557 564 566 568 569

571 571 572 573 573 574 574 575 575

Contents

Switchable Hydrophilic/Hydrophobic Fluorocarbon Layer Obtained on Porous Alumina using c-C4F8 Plasma Treatment 19.5 Comparison of Surface Modifications of Porous Alumina using Various Fluorinated Media: CF4, C3F8 and c-C4F8 19.6 Conclusion Acknowledgements References

xv

19.4.2

Index

578 580 581 582 582 583

Preface

Fluorides and fluorinated materials affect various aspects of modern life. The strategic importance of fluoride materials, and the use of adapted fluorination surface treatments, concern many research fields and applications in areas such as energy production, microelectronics and photonics, catalysis, colour pigments, textiles, cosmetics, plastics, domestic wares, automotive technology and building. Among the issues with which they are concerned [1–4] are: • the historical importance of fluoride fluxes in the production of metals, in particular aluminium; • the critical place of fluorine and fluorides in conversion energy processes – for example components of Li-ion batteries and fuel cells, enrichment of 235U through uranium hexafluoride for nuclear energy; • the etching of silicon wafers for microelectronics; • the technical revolution of fluoropolymers and fluoride coatings, for example TeflonÒ and fluorinated plastics, waterproof clothes, biomaterials for cardiovascular or retinal surgeries, kitchen wares, and so forth; • the beneficial influence of fluoride on dental caries; • the dominant use of fluorinated molecules in agrochemistry and phytosanitary products; • the dramatic increase of fluorine-containing molecules for medicine and pharmacy, as efficient imaging products, as dental composites for cariostatic improvement, and so forth; • the use of 18F-labelled molecules in positron emission tomography (PET) for early diagnosis of cancer and Alzheimer’s disease. In the case of inorganic fluorinated solids, numerous improvements have recently been achieved through the elaboration and functionalization of the materials on a nanometric scale. The present book covers several classes of nanostructured and functionalized inorganic fluorides, oxide-fluorides, hybrids, mesoporous materials and fluorinated oxides such as silica and alumina. The morphologies concerned range from powders or glassceramics to thin layers and coatings whereas the applications involved include catalysts, inorganic charges, superconductors, ionic conductors, ultaviolet (UV) absorbers, phosphors, materials for integrated optics, and so forth. Several books have been devoted to the reactivity of carbon-based materials with fluorine (carbon fibres, fullerene, carbon nanotubes, etc) [1,2,5,6], so these types of materials will not be treated in the present book.

xviii

Preface

The book arose from discussions that took place during the FUNFLUOS project (2004–2008), carried out within the Sixth European Framework Programme. This project involved about ten groups from Germany, France, Slovenia and the UK, all aimed at the synthesis and characterization of fluorinated materials with properties tailored for specific applications. The topics appearing in the book range from new synthesis routes to physical-chemical characterizations. They address important properties of these materials, including morphology, structure, thermal stability, superconductivity, magnetism, spectroscopic and optical behaviour. Detailed ab initio investigations and simulations provide a comparison with experimental results, and potential applications of the final products are also proposed. In the first section, two innovative routes toward nanoscaled metal fluorides and hydroxyfluorides are presented: the fluorolytic sol-gel synthesis by E. Kemnitz et al. and the microwave-assisted route by D. Dambournet et al. In a second section, several physical-chemical characterizations are developed in order to understand the mechanisms that are responsible for the improvement of the properties of these materials: investigation of the main characteristics of high-surface-area aluminium fluorides as catalysts by E. Kemnitz and S. Ru¨diger; determination of surface acidities (Lewis and Brønsted types) using a large range of probe molecules, by A. Vimont et al.; a better knowledge of the environment of the different nuclei using high-resolution solid-state nuclear magnetic resonance (NMR) by C. Legein et al. The theoretical investigation of these topics is highlighted by the predictive modelling of aluminium fluoride surfaces by C. Bailey et al., which allows a better understanding of the underlying processes at the molecular and nano levels. An example of industrial application of the inorganic fluorides is given by P. Garcia Juan et al. In the following section, some examples of outstanding optical properties of nanostructured fluorides are proposed: nanostructured fluorocompounds as UV absorbers, by A. Demourgues et al.; transparent oxyfluoride glass-ceramics by M. Mortier and G. Dantelle; luminescent and antireflective coating of (oxy)fluorinated materials obtained by the sol-gel technique, by S. Fujihara; planar optical waveguides based on fluoride glasses, by B. Boulard. Hybrids, composites and mesoporous fluorides are original materials with great potential and the interesting nature of such materials is illustrated in the next section by the chapters on polyanion condensation in inorganic-organic hybrid fluorides, by K. Adil et al.; superconducting/magnetic properties of Cuand Mn-based oxyfluorides, by E. Antipov and A. Abakumov; ionic conductivity of fluoride-containing phases by E. Ardashnikova et al.; intercalation in hybrid compounds containing perfluoroalkyl groups, by Y. Matsuo. The two following chapters deal with the synthesis of microporous frameworks using the fluoride and F2-gas routes, respectively. The examples concern either compounds based on silica, germanium, phosphates and clays, by J. L. Paillaud et al., or highly fluorinated silica, by A. Demourgues et al. The optical and magnetic properties of oxyfluoride glasses based on rare-earth elements are illustrated by S. Yonezawa et al. Finally the chapter by A. Tressaud et al. describes the use of surface fluorination of porous alumina for applications in offset technology. A very wide range of materials, properties, and applications have therefore been gathered in this book, which covers various new fields in which inorganic fluorides are part of the innovating process. Among the information that can bring answers to some crucial questions in materials science, we can quote new synthesis routes towards more

Preface

xix

efficient and less aggressive catalysts, protection against harmful UV radiation, new integrated lasers and optical amplifiers, antireflective coatings, solid-state ionic conductors, highly hydrophobic silica and switchable coatings for offset technology. Erhard Kemnitz and Alain Tressaud Berlin and Bordeaux September 2009

References [1] Advanced Inorganic Fluorides, T. Nakajima, B. Zemva, A. Tressaud (Eds), Elsevier, Amsterdam (2000). [2] Fluorinated Materials for Energy Storage, T. Nakajima, H. Groult (Eds), Elsevier, Amsterdam (2005). [3] Fluorine and the Environment, Vol. 1 and Vol. 2, A. Tressaud (Ed.), Elsevier, Amsterdam (2006). [4] Fluorine and Health, A. Tressaud and G. Haufe (Eds), Elsevier, Amsterdam (2008). [5] Graphite Fluorides and Carbon-Fluorine Compounds, T. Nakajima (Ed.), CRC Press, Boca Raton, FL (1991). [6] ‘Fluorofullerenes’, in Dekker Encyclopedia of Nanoscience and Nanotechnology, O. V. Boltalina, S. H. Strauss, 2nd edition, Dekker, New York (2009).

The Funfluos European Network (2004): First row (from left to right): D. Menz, B. Zˇemva, E. Kemnitz (Coordinator), A. Demourgues, A. Tressaud, and J. Winfield. Second row (from left to right): U. Gross, M. Feist (partly hidden), S. Ru¨diger, P. Millet (European Commission), N. Harrison, A. Wander, T. Skapin and S. Schro¨der

List of Contributors Artem M. Abakumov, Department of Chemistry, Moscow State, University, Moscow, Russia Karim Adil, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France Evgeny V. Antipov, Department of Chemistry, Moscow State, University, Moscow, Russia Elena I. Ardashnikova, Department of Chemistry, Moscow State, University, Moscow, Russia Christine L. Bailey, Computational Science and Engineering Department, STFC Daresbury Laboratory, Warrington, Cheshire, UK Monique Body, Laboratoire de Physique de l’Etat Condense´, UMR-CNRS, Universite´ der Maine, Le Mans, France Brigitte Boulard, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France Jocelyne Brendle´, Laboratoire de Mate´riaux a` Porosite´ Controˆle´e, UMR-CNRS, Universite´ de Haute Alsace, Mulhouse, France Jean-Yves Buzare´, Laboratoire de Physique de l’Etat Condense´, UMR-CNRS, Universite´ der Maine, Le Mans, France Amandine Cadiau, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France Philippe Caullet, Laboratoire de Mate´riaux a` Porosite´ Controˆle´e, UMR-CNRS, Universite´ de Haute Alsace, Mulhouse, France Damien Dambournet, Institute of Condensed Matter Chemistry of Bordeaux (ICMCBCNRS), University Bordeaux 1, Pessac, France Ge´raldine Dantelle, Laboratoire de Photonique Quantique et Mole´culaire (LPQM), UMR CNRS, Cachan, France

xxii

List of Contributors

Marco Daturi, ENSICAEN, Universite´ de Caen, CNRS, Caen, France Alain Demourgues, Institute of Condensed Matter Chemistry of Bordeaux (ICMCBCNRS), University Bordeaux 1, Pessac, France Etienne Durand, Institute of Condensed Matter Chemistry of Bordeaux (ICMCB-CNRS), University Bordeaux 1, Pessac, France Johannes Eicher, Solvay Fluor GmbH, Hannover, Germany Shinobu Fujihara, Department of Applied Chemistry, Faculty of Science and Technology, Keio University, Yokohama, Japan Placido Garcia Juan, Solvay Fluor GmbH, Hannover, Germany Nicholas Harrison, Computational Science and Engineering Department, STFC Daresbury, Laboratory, Warrington, Cheshire, UK Department of Chemistry, Imperial College London, London, UK Annie He´mon-Ribaud, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France Erhard Kemnitz, Institute for Chemistry, Humboldt University of Berlin, Berlin, Germany Jae-ho Kim, Graduate School of Engineering, University of Fukui, Fukui, Japan Ilya B. Kutsenok, Department of Chemistry, Moscow State, University, Moscow, Russia Christine Labruge`re, Institute of Condensed Matter Chemistry of Bordeaux (ICMCBCNRS), University Bordeaux 1, Pessac, France Emilie Lataste, Institute of Condensed Matter Chemistry of Bordeaux (ICMCB-CNRS), University Bordeaux 1, Pessac, France Marc Leblanc, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France Christophe Legein, Laboratoire des Oxydes et Fluorures, CNRS, Le Mans, France Vincent Maisonneuve, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France Charlotte Martineau, Laboratoire des Oxydes et Fluorures, UMR CNRS, Le Mans, France, Tectospin, Universite´ de Versailles Saint Quentin en Yvelines, Versailles, France Yoshiaki Matsuo, Department of Materials Science and Chemistry, University of Hyogo, Hyogo, Japan

List of Contributors

xxiii

Michel Mortier, Laboratoire de Chimie de la Matie`re Condense´e de Paris, UMR CNRS, Paris, France Sanghamitra Mukhopadhyay, Department of Chemistry, Imperial College London, London, UK Jean-Louis Paillaud, Laboratoire de Mate´riaux a` Porosite´ Controˆle´e, UMR-CNRS, Universite´ de Haute Alsace, Mulhouse, France Joe¨l Patarin, Laboratoire de Mate´riaux a` Porosite´ Controˆle´e, UMR-CNRS, Universite´ de Haute Alsace, Mulhouse, France Nicolas Penin, Institute of Condensed Matter Chemistry of Bordeaux (ICMCB-CNRS), University Bordeaux 1, Pessac, France Vladimir A. Prituzhalov, Department of Chemistry, Moscow State, University, Moscow, Russia Stephan Ru¨diger, Institute for Chemistry, Humboldt University of Berlin, Berlin, Germany Gudrun Scholz, Institute for Chemistry, Humboldt University of Berlin, Berlin, Germany Thomas Schwarze, Solvay Fluor GmbH, Hannover, Germany Barry Searle, Computational Science and Engineering Department, STFC Daresbury Laboratory, Warrington, Cheshire, UK Gilles Silly, Institut Charles Gerhardt Montpellier, Physicochimie des Mate´riaux De´sordonne´s et Poreux, Universite´ de Montpellier II, Montpellier, France Ange´lique Simon-Masseron, Laboratoire de Mate´riaux a` Porosite´ Controˆle´e, UMRCNRS, Universite´ de Haute Alsace, Mulhouse, France Laetitia Sronek, Institute of Condensed Matter Chemistry of Bordeaux (ICMCB-CNRS), University Bordeaux 1, Pessac, France Hans-Walter Swidersky, Solvay Fluor GmbH, Hannover, Germany Masayuki Takashima, Graduate School of Engineering, University of Fukui, Fukui, Japan Alain Tressaud, Institute of Condensed Matter Chemistry of Bordeaux (ICMCB-CNRS), University Bordeaux 1, Pessac, France

xxiv

List of Contributors

Alexandre Vimont, ENSICAEN, Universite´ de Caen, CNRS, Caen, France Adrian Wander, Computational Science and Engineering Department, STFC Daresbury Laboratory, Warrington, Cheshire, UK John M. Winfield, Department of Chemistry, University of Glasgow, Glasgow, UK Susumu Yonezawa, Graduate School of Engineering, University of Fukui, Fukui, Japan

1 Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties Erhard Kemnitz, Gudrun Scholz and Stephan Ru¨diger Humboldt-Universita¨t zu Berlin, Institut fu¨r Chemie, Brook – Taylor – Str. 2, D – 12489 Berlin, Germany

1.1

Introduction

Sols are dispersions of nanoscopic solid particles in, for example, liquids – i.e., colloidal solutions. The particles can agglomerate forming a three-dimensional network in the presence of large amounts of the liquid thus forming a gel. Inorganic sols are prepared via the sol-gel process, the investigation of which started in the nineteenth century. This process received great impetus from the investigations of Sto¨ber et al. [1], who studied the use of pH adjustment on the size of silica particles prepared via sol-gel hydrolysis of tetraalkoxysilanes. 1.1.1

Sol-Gel Syntheses of Oxides – An Intensively Studied and Widely Used Process

Hydrolysis of alkoxysilanes and later on of metal alkoxides in organic solutions has become an intensively studied and widely used process [2]. The most common products are almost homodispersed nanosized silica or metal oxide particles for, e.g., ceramics or

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Edited by Alain Tressaud 2010 John Wiley & Sons, Ltd

2

Functionalized Inorganic Fluorides

glasses, or the aqueous sols are used to prepare different coatings for, e.g., optical purposes. Optical applications depend on differences in the respective indices of refraction of the coated material and the applied layer. The latter has to be of very uniform and thoroughly adjusted thickness. The sol-gel hydrolysis of alkoxysilanes, the most intensively explored one, basically proceeds in two steps. The first step is the hydrolytic replacement of alkoxy groups, OR, by hydroxyl groups, OH, shown schematically in Equation (1.1) for the first alkoxy group: SiðORÞ4 þ H2 O ! ðRO Þ 3 SiOH þ ROH

(1:1)

Because of their relatively high hydrolytic stability, hydrolysis of alkoxysilanes (1.1) has to be catalysed by Brønstedt acids or bases. In the second step, the primary hydrolysis products undergo condensation reactions under elimination of water (Equation (1.2)) or alcohol (Equation (1.3)). X3 Si-OH þ HO-SiX3 ! X3 Si-O-SiX3 þ H2 OðX ¼ OR; OHÞ

(1:2)

X3 Si-OH þ RO-SiX3 ! X3 Si-O-SiX3 þ ROHðX ¼ OR; OHÞ

(1:3)

As a result tiny particles with a very open structure are formed. The overall process can be controlled by adjusting the reaction conditions. The colloidal solution of these particles, the sol, can be used as such for, e.g., coating or it can be worked up to yield, eventually, nanoscopic oxide particles. However, metal oxide sols obtained in this way always contain a sometimes remarkable organic part. Its separation demands calcination temperatures of at least 623 K in order to convert the ‘precursors’ into pure metal oxide materials. Substituting a certain part of the alkoxidic groups by nonhydrolysable ones, such as alkyl groups in the case of alkoxysilanes or phosphonic acid in the case of metal alkoxides, organically modified oxides, i.e. inorganic-organic hybrid materials, have been prepared.

1.1.2

Sol-Gel Syntheses of Metal Fluorides – Overview of Methods

Selected metal fluorides can, in application-relevant fields, outperform metal oxides and silica. Thus, for instance, magnesium fluoride and aluminium fluoride and, in particular, alkali hexafluoroaluminates have both a lower index of refraction and a much broader spectral range of transparency even than silica, making them very interesting for optical layers. Consequently, several approaches for the preparation of nanoscopic metal fluorides and metal fluoride thin layers have been developed and proposed. Besides physical methods such as milling, laser dispersion or molecular-beam epitaxy, different chemical methods exist. Basically, three approaches can be distinguished: (i) Postfluorination of a metal oxide preformed via sol-gel route [3]. This route is shown schematically in Scheme 1.1. The disadvantages of this route are, to name two, incomplete fluorination of the bulk metal oxides and decrease of surface area in course of the fluorination.

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties M(OR)x + xH2O hydrolysis

MOx/2(gel) bulk

3

MOx/2 (sol) + xROH coating

MOx/2(gel) film fluorination

fluorination + calcination MFx or MOx/2/MFx glass

MFx or MOx/2 /MFx film

Scheme 1.1 Metal fluoride preparation via post fluorination of sol-gel prepared metal oxides (Reproduced from [4] by permission of Elsevier Publishers)

(ii) Preparation of a precursor containing a metal compound with organically bound fluorine such as trifluoroacetate, which is calcined to decompose the fluoroorganic component under formation of metal fluoride [5]. This route, shown in Scheme 1.2, also starts from metal alkoxides, which are reacted in solution with, e.g., trifluoroacetic acid to form metal trifluoroacetate sol. This can be M(OR)x and/or MOAc + CF3COOH + H2O + org. solvent

metal trifluoroacetate (TFA) precursor solution drying

metal-TFA(gel) bulk

coating

metal-TFA(gel) film heating

heating metal fluoride powder

metal fluoride film

heating at higher temperature

MOx/2 /MFx powder

MOx/2 /MFx film

Scheme 1.2 Metal fluoride preparation via metal fluoroacetate sol-gel formation and following thermal decomposition. (Reproduced from [4] by permission of Elsevier Publishers)

4

Functionalized Inorganic Fluorides

used for coating experiments. The decisive final step is the thermal decomposition of the fluoro-organic constituent, because of which thermolabile materials cannot be coated. Another disadvantage is the probability that oxidic components can be formed as admixtures or oxofluorides. (iii) Fluorolytic sol-gel process as counterpart to the hydrolytic one. The fluorolytic sol-gel route follows rather strictly the ‘classical’ hydrolytic one by reacting metal alkoxides in anhydrous solution with hydrogen fluoride instead of the hydrogen oxide of the ‘classical’ process. Consequently, it results eventually in metal fluorides instead of metal oxides. The fluorolytic sol-gel process, its execution, mechanism, scope as well as properties and possible fields of application of its products are the subjects of this chapter.

1.2

Fluorolytic Sol-Gel Synthesis

Metal alkoxides can be regarded as metal salts of alcohols, where the latter are very weak Brønstedt acids. Acids that are stronger than the respective alcohol can therefore replace alkoxy groups attached to the metal ion under liberation of the alcohol and formation of the metal fluoride according to Equation (1.4). MðORÞn þ x HF ! MðORÞn x Fx þ x ROH ðM ¼ metal ionÞ

(1:4)

In fact, starting with aluminium isopropoxide [6], a broad range of metal alkoxides have been subjected to a sol-gel-like liquid-phase fluorination with hydrogen fluoride in organic solution [4, 7]. Although Equation (1.4) closely resembled Equation (1.1) there is an important difference in that condensation reactions like those of Equations (1.2) and (1.3) are not possible in the fluorolysis system. On the other hand, the fluorolysis reactions typically result in the formation of a sol-gel. The formation of a gel was already mentioned in the first paper on metal alkoxide fluorolysis, reporting the reaction of aluminium isopropoxide in alcoholic solution with an ethereal solution of hydrogen fluoride [6]. The gel formation is obviously due to an important consequence of the replacement of alkoxy groups by fluoride, i.e., the Lewis acidity of the metal ion increases leading to a strengthening of the interaction between (liberated) alcohol molecules and metal ions. As a result alcohol molecules that can occupy ligand positions might establish a loose net between (partly) fluorinated metal ions resulting eventually in metal fluoride sol or even gels. Surprisingly, attempts to isolate pure AlF3 by drying and calcining the gel were not successful; the product obtained had an understoichiometric amount of fluorine even when the primary reaction has been carried out with an overstoichiometric amount of HF [8]. An additional fluorination of the dried gel under gentle conditions (see below) has proved to be a suitable way to remove the attached organic components resulting in X-ray amorphous, highly Lewis acidic aluminium fluoride with unusual large specific surface area, named HS-AlF3 [9].

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

1.2.1

5

Mechanism and Properties

To gain insight into the fluorolytic sol-gel process, elucidating the mechanism and optimizing both the experimental procedure and the properties of the products, the influence of all adjustable synthesis parameters was tested as well as the process. In particular, the products were analysed by a broad range of analytical, experimental and theoretical methods. Most of these investigations focussed on the synthesis and properties of HS-AlF3, followed by HS-MgF2. The results and consequences of these investigations are presented and discussed below as well as in subsequent chapters. 1.2.1.1

Approach to Mechanism and Resulting Properties

The fluorolytic sol-gel synthesis of metal fluorides is summarized in Scheme 1.3. [M(OR)n]x

+S

[M(OR)n]y /S

+

HF/S’

1 2 M-F-SS’ sol 3 M-F-SS’ wet gel 4 M-F dry gel 5 HS-MFn

Scheme 1.3 Fluorolytic sol-gel synthesis of high surface area metal fluorides (for explanations of the numbers 1 to 5 see text)

The synthesis was investigated in detail, aiming its optimization for HS-AlF3 [8] and also for HS-MgF2 [10, 11]. In short, the metal alkoxide, the structure of which can be very complex [12], is dissolved in an alcohol or another suitable organic solvent (step 1 in Scheme 1.3), so that in the case of aluminium isopropoxide the tetrameric structure is predominantly preserved. A solution of hydrogen fluoride in, e.g., ether or alcohol is added in approximately stoichiometric amounts to the alkoxide solution. Ratios of HF:Al from 2 to 4 are well tolerated, resulting in a clear, translucent sol (step 2 in Scheme 1.3), which more-or-less rapidly becomes a gel, depending on concentration and type of solvent (step 3 in Scheme 1.3). Varying the type of alkoxide, from methoxide to ethoxide, isopropoxide to butoxide, did not markedly affect the outcome of the fluorolytic reaction.The only meaningful criteria for estimation of these and other synthesis parameters were thus the surface area and especially Lewis acidity of the final HS-AlF3.

6

Functionalized Inorganic Fluorides

There is yet another route to the synthesis of metal fluoride sol-gel, exemplified for aluminium and magnesium, namely the direct reaction of the metal with an alcoholic HF solution [13]. Upon drying the sol-gel under mild conditions, at about 343 K under vacuum or freeze drying, or under microwave irradiation, a solid, X-ray amorphous dry gel is formed (step 4 in Scheme 1.3), which contains, in case of the Al-F-system, large amounts of organic material, indicated by a carbon content of about 20 %–30 %. An empirical formula based on elemental analysis for a dry gel prepared from Al(OiPr)3 in iPrOH is AlF2.7[OCH(CH3)2]0.30.7-0.8(CH3)2CHOH. With metal ions of lower Lewis acidity, decisive lower carbon contents were found, e.g. 3 %–7 % C in the Mg-F-system. Obviously, part of the alcohol, which is the predominant constituent of the wet gel, is very tightly attached to the highly Lewis acidic Al3þ ion, as can also be seen from its thermo-analysis (Figure 1.1). The weakly endothermic mass loss of about 24 % up to 473 K can clearly be attributed to the release of solvating iPrOH and the more pronounced smaller one around 495 K to the split off of alkoxide groups. Evaluating thermal analysis data of many different experiments, it became obvious that the mass loss proceeds stepwise. The alcohol content of the steps corresponds to the respective compositions of about AlF3:1ROH after heating at 343 K under vacuum, about AlF3:0.45ROH after heating at 573 K in N2, and about AlF3:0.1ROH after continued heating up to about 600 K [13]. The exothermal peak in Figure 1.1 at 836 K is due to crystallization, i.e. of -AlF3 formation. In order to obtain a still X-ray amorphous aluminium fluoride, the dry gel with its understoichiometric fluorine content has to be freed from its organic constituents under fluorinating conditions. This can be accomplished in a gas-solid reaction at elevated temperatures up to 573 K with vaporized fluorocarbon compounds such as CHClF2 diluted with an inert gas. An aluminium fluoride is obtained, named HS-AlF3, which is still amorphous, has a specific surface area of about 200 m2/g and shows extremely high Lewis acidity (see below). An unwanted consequence of the extreme Lewis acidity is the readily occurring coke formation preventing the use of fluorochloroethane compounds as fluorinating agents. The postfluorination step is essential for HS-AlF3 formation, therefore all parameters have been comprehensively tested, such as type and concentration of the fluorinating agent, flow rate, temperature, and also ageing of the Al-F-sol. Even under optimum conditions with CHClF2 or CH2F2 the formation of black spots or sometimes of ‘channels’ could be observed indicating that at these spots the formation of HS-AlF3 had started, which then subsequently catalysed coke formation. Such side reactions are, for obvious reasons, not possible using HF as fluorinating agent. In addition, the stronger fluorinating HF can be used at lower temperatures and should preserve the amorphous state with its high surface area even better. The latter criterion could only be fulfilled using rather diluted HF, obviously to reduce the otherwise high reaction enthalpy. However, ‘HS-AlF3’ prepared with gaseous HF did not show the expected Lewis acidity. It turned out that HF had behaved as base, which became attached to the strongest Lewis acid sites of the solid, thereby blocking them. Only by additional longer flushing with a stream of inert gas or, even better, of CHClF2 vapour at elevated temperatures (up to 573 K) did the material become the expected strong solid Lewis acid, i.e. HS-AlF3. HS-AlF3 could also be prepared with elemental fluorine using a dry Al-F-gel of low carbon content obtained by microwave heating of an Al-F-sol under autologous pressure followed by microwave-assisted drying.

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties DTA /μV DTG /(%/min)

TG (%) TG 100 90 80

–23.85%

50

10

DTG

a –9.48%

70 DTA 60

–5.65%

↑exo

472 836

415

821

30 373

TG (%) TG 100

473

573 673 Temperature/K

773

DTG /(%/min) Ion Current –11∗10/A [2]

m41

80 DTG

b

70 60 m87 (x100)

30 473

573 673 Temperature/K

–5

–15

m45

373

0

0 –0.50 –1.00 –1.50 –2.00 –2.50 –3.00 –3.50

873

90

40

5

–10

495

40

50

7

773

0 –0.50 –1.00 –1.50 –2.00 –2.50 –3.00 –3.50

6.0 5.0 4.0 3.0 2.0 1.0 0

873

Figure 1.1 TA-MS of AlF2.7[OCH(CH3)2]0.30.70.8(CH3)2CHOH; (a) Thermoanalytical curves in N2 (19.50 mg); (b) Ion current curves for m/z 41 (C3H5þ), m/z 45 (C2H5Oþ), and m/z 87 (C5H11Oþ), indicating the release of propene, i-propanol, and diisopropyl ether, respectively. (Reprinted with permission from [9] Copyright (2005) RSC.)

The results of experimental investigation are in accordance with the following tentative mechanism of the sol-gel fluorination. Upon addition of HF to the alkoxide solution a stepwise replacement of –OR by –F starts, whereby the coordinating alcohol as linking group prevents the formation of a three dimensional purely F-linked crystal. The sol-gel state, almost immediately formed, kinetically prevents the stoichiometric fluorination of the aluminium species. Upon removing the alcohol under gentle conditions the gel structure only partly collapses; alcohol molecules of the first ligand sphere remain obviously attached to Al. When these molecules are removed under appropriate mild conditions there is no crystallization taking place but the disordered X-ray amorphous state connected with an unusual high surface area remains and a part of the Al atoms becomes co-ordinately unsaturated consequently exhibiting very high Lewis acidity (see below). For HS-AlF3 a surface area up to 400 m2/g has been determined by N2 adsorption/desorption experiments. Typical isotherms are shown in Figure 1.2.

Functionalized Inorganic Fluorides Pore volume dV/d logD [cm3/g]

8

Volume adsorbed [cm3/g]

250

200

150

2

1

0 0

100 200 300 Pore diameter [Å]

100

50 adsorption desorption 0 0.0

0.1

0.2

0.3 0.4 0.5 0.6 0.7 Relative Pressure (p/p 0)

0.8

0.9

1.0

Figure 1.2 N2 adsorption/desorption isotherms and pore size distribution of HS-AlF3 (Reprinted with permission from [6] Copyright (2003) Wiley-VCH.)

The tentative mechanism is given in more detail and supported by the analytical investigations discussed in the following section.

1.2.2

Insight into Mechanism by Analytical Methods

Different spectroscopic, microscopic and diffraction methods like IR and Raman spectroscopy, TEM or XRD were applied to characterize educts, intermediates or products of the sol-gel process. For a detailed insight into the mechanism of the fluorolytic sol-gel process, however, the application of NMR spectroscopy is the method of choice. The NMR experiments, both in liquid and in solid state, allow direct observations of local structures and their changes even if the matrices suffer from a loss of lattice periodicity. For the fluorolytic so-gel process both liquid state NMR experiments were realized for the alkoxide solutions, sols and thin gels as well as solid state MAS NMR experiments for the alkoxide, dried alkoxide fluoride gels and high-surface fluorides including 1H, 13C, 27 Al and 19F as sensitive spin probes. The same spin probes along with the use of 1D and 2D NMR experiments allow local structures in liquids and solids to be addressed and directly compared, to follow their changes with a progressive degree of fluorination and finally to derive a possible reaction pathway for this process. Due to the good experimental accessibility of the mentioned spin probes, a detailed study was conducted for the reaction steps ending up with HS-AlF3. Results of these studies are presented briefly in the following section.

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

1.2.2.1

9

Sol-Gel Reaction

1.2.2.1.1 Aluminiumisopropoxide (Al(OiPr)3) as Solid Precursor Compound The basis for an understanding of changes during the fluorination process is the knowledge and assignment of molecular structures existing in solid Al(OiPr)3 and its solution in iPrOH. 27 Al MAS NMR measurements of solid Al(OiPr)3 give unambiguous indications for the existence of two distinguishable aluminium sites in the matrix in agreement with XRD findings [14, 15] (Figure 1.3). A simulation of the spectrum (Figure 1.3) results in typical chemical shift values for AlO4 (di ¼ 61.5 ppm) and AlO6 (di ¼ 1.7 ppm) units [16, 17]. Both the kind and the intensity ratio of AlO4 : AlO6 as 3 : 1 confirm the tetrameric molecular structure (see also Figure 1.5, 1) existing in the tetragonal crystal structure [14, 15].

50

400

200

0

0

(ppm)

–200

–50

–100

–400

(kHz)

Figure 1.3 27Al MAS NMR spectrum of Al(OiPr)3 ( r ¼ 20 kHz; B0 ¼ 9.4 T) (solid line: experiment, dotted line: simulation; insert: central transitions with quadrupolar splitting) (Reprinted with permission from [16] Copyright (2006) Humboldt University.)

In addition to 27Al, the respective 1H-13C CP MAS NMR spectrum allows a distinction to be made not only between CHO and CH3 groups but also between terminal (Al-O) and bridging (Al-O-Al) isopropoxide units. Very narrow line widths in part allow the assignment of 18 distinguishable carbon sites in the NMR spectrum in agreement with the crystal structure (see Figure 1.10 (top)) [14, 15]. According to suggestions made by Abraham [17], terminal (63 ppm) and bridged (66 ppm) CHO – groups are located in the low-field part of the spectrum; terminal CH3 groups are detected in the range between 28–30 ppm and bridging CH3 groups dominate the high-field part of the spectrum (see Figure 1.10 (top)), [16]). 1.2.2.1.2 Aluminiumisopropoxide (Al(OiPr)3) Dissolved in Isopropanol First 1H NMR studies on possible structures of Al(OiPr)3 in different solutions range back to 1963 [18] followed by first 27Al NMR measurements in 1973 [19]. Since that

Functionalized Inorganic Fluorides

10

time aluminium isopropoxide solutions and possible species therein, using CCl4, benzene or toluene as solvents, are subjects of current interest applying different NMR techniques [20–22]. Isopropanol is the standard solvent used for the fluorolytic sol-gel synthesis [4, 6–8], so results of a careful reinvestigation of possible isopropanolic solution are presented in summary. 27 Al NMR spectra of Al(OiPr)3 dissolved in isopropanol are shown in Figure 1.4a,b; the spectrum recorded for the same sample dissolved in diethylether is depicted in Figure 1.4c.

a

b

250

200

150

100 50 (ppm)

0

–50 –100

c

250

200

150

100 50 (ppm)

0

–50 –100

Figure 1.4 27Al NMR spectra of different Al(OiPr)3 – solutions: a and b Al(OiPr)3 in iPrOH recorded with a 400 MHz – spectrometer (a) and a 600 MHz – spectrometer (b), c Al(OiPr)3 in diethylether (400 MHz – spectrometer). For all: solid: experimental spectrum, dashed: simulation and dotted: decomposition. (Reprinted with permission from [23] Copyright (2007) American Chemical Society.)

The spectrum obtained for the etheric solution (Figure 1.4c) indicates the existence of only tetrameric aluminium isopropoxide species in solution (AlO6 (2.5 ppm, 26.7 %), AlO4 (61.8 ppm, 73.3 %), 1 in Figure 1.5) but the situation in isopropanolic solution is much more complex. Applying two different fields (Figure 1.4a,b), both spectra reveal, beside AlO6 and AlO4, the existence of an additional fivefold coordinated aluminium species AlO5 at 32 ppm. They support therewith the existence of trimeric Al(OiPr)3 species (3 in Figure 1.5). Calculated intensities of the spectra support the assumption of further species. Cyclic trimeric species (2 in Figure 1.5) with possible distorted AlO4 polyhedra explain the strong low-field shifted resonance observed at 85 ppm [16, 23].

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

11

O O Al O

Al O

O

O O Al O O O Al O O 1

O

O O O Al O Al O O O Al O O

Al O O O O Al Al O O O

2

3

O

O = OiPr

Figure 1.5 Possible structural units of Al(OiPr)3 in solution. (Reprinted with permission from [23] Copyright (2007) American Chemical Society.)

Moreover, an equilibrium between different trimeric species seems to be feasible, which involves a trimeric species with central sixfold oxygen-coordinated aluminium. The latter may result from the interaction of trimeric species 3 (Figure 1.5) and a solvent molecule as shown in Figure 1.6 [16].

O Al

O Al O

O O O

Al

O

HOiPr

O O

HOiPr

Al O

O O

Al O

O

O O

Al O

O 4

4'

O : OiPr

Figure 1.6 Possible equilibria between different trimeric Al(OiPr)3 species in isopropanolic solution. (Reprinted with permission from [16] Copyright (2006) Humboldt University.)

In the appropriate 1H and 13C NMR spectra, no indications are observable that distinguish between tetramers 1 and trimers 2 and 3 (Figure 1.5), respectively. The only possible discrimination between bridging and terminal isopropoxide groups holds for all of the mentioned Al(OiPr)3 species (Figure 1.5) [16]. Summing up these results it can be concluded that the isopropanolic solution of Al(OiPr)3 is dominated by the tetrameric Al(OiPr)3 species 1 (Figure 1.5) accompanied by a smaller amount of trimeric species.

1.2.2.2

Structural Changes at the Fluorination Process in Sols and Thin Gels

On the basis of the 27Al NMR spectra of isopropanolic Al(OiPr)3 solution, structural changes were followed after adding increasing amounts of HF/iPrOH solution to the aluminium isopropoxide solution according to Equation (1.4). In dependence on the molar Al:F ratio the appearance of the reaction mixture ranges from clear sols (4:1) to opaque gels (1:3). Figure 1.7 represents the 27Al and 19F NMR spectra obtained with rising content of fluorine (from a to d). It is obvious that in this order the amount of AlO6 species

Functionalized Inorganic Fluorides

12

(narrow line in Figure 1.4a,b; central units of 1 (Figure 1.5) is decreased. In the same manner a decreasing proportion of the sum over all fourfold Al species is found and, contrary to that, an increasing amount of fivefold coordinated aluminium species AlO5 (signals at about 35 ppm) [23].

27Al

19

F

–156 ppm –147 ppm

a

a –165 ppm –163 ppm

b

c

d

–161 ppm

b

–171 ppm

c

d

500 400 300 200 100 0 –100 –200 –300 –400 –100 (ppm)

–120

–140 –160 (ppm)

–180

–200

Figure 1.7 27Al NMR and 19F NMR spectra of different sols and wet gels(B0¼9.4 T). For all: solid line: experimental spectrum, dashed: simulation, dotted: decomposition. From a to d increasing content of fluorine. Molar ratios Al: F: (a) 4 : 1, (b) 2 : 1, (c) 1 : 1, (d) 1 : 2. (Reprinted with permission from [23] Copyright (2007) American Chemical Society.)

All 19F spectra are characterized by a group of three sharp signals (–161 ppm, –163 ppm, –165 ppm) with different intensities. All these signals are in a typical region for fluorine bounded on aluminium centres in a mixed oxygen-fluorine coordination with different fluorine ratios [13, 24–28]. With higher fluorine content, the intensity of the 19F NMR spectrum is more and more dominated by a broad peak at about –160 ppm (Figure 1.7, 19 F,d). These line-broadening effects result mainly from 19F-19F homonuclear dipolar couplings ending up in one broad peak in the static 19F NMR spectrum for the gel with molar ratio Al:F as 1:3. 1H and 13C NMR spectra of sols and gels show two main effects

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

13

with increasing fluorine supply: the portion of bridging isopropoxide groups is decreasing while the portion of terminal isopropoxide groups is affected only with higher fluorine supply [16, 23]. Obviously, the fluorination starts by protonation of bridging isopropoxide groups with the consequence of a line broadening of the corresponding signals, while the intensities of signals of terminal groups first remain constant. This first step is also supported by recently accomplished DFT calculations [23, 29]. A subsequent fluorination step may involve an attack of fluorine ions or HF on the central Al atoms of 1 (Figure 1.5) and the substitution of the protonated isopropoxide group by fluorine. 1.2.2.2.1 Isolated Single Crystals as Intermediates Evidence for a stepwise introduction of fluorine into the coordination sphere of Al is also given by the successful isolation of single crystals of partly fluorinated aluminium as well as magnesium alkoxide fluorides, Al3(OiPr)8FDMSO (Figure 1.8a) [9], Al3(OiPr)8FPy (Figure 1.8b) [30] and Mg6F2(OMe) 10(MeOH)14 (Figure 1.11) [31, 32]. In each case the isolation of single crystals works only if the fluorine supply is very low, i.e. Al:F or Mg:F > 1.

O(6) O(8)

O(1)

Al(3)

Al2

Al(2)

O(3)

O(5)

O(2) O(7)

O6

O5

O1

O3

O8

O2

O(4) Al(1)

Al3

Al1

O4

O(9) O7

N1 F1

F S

a

b

Figure 1.8 Crystal structure of (a) Al3(OiPr)8FDMSO (Reprinted with permission from [9] Copyright (2005) RSC); (b) Al3(OiPr)8FPy (Reprinted with permission from [30] Copyright (2008) Humboldt University of Berlin.)

Three distinguishable aluminium sites and one fluorine site are expected for Al3(OiPr)8FPy in the 27Al and 19F MAS NMR spectra, which are given in Figure 1.9. Simulation of the 27Al NMR spectrum supports this assumption and the decomposition obtained is given in Figure 1.9. The chemical shift values are close to those of crystalline Al(OiPr)3 emphasizing their structural similarity. As a consequence of the coordination of fluorine and a solvent molecule the originally higher symmetric AlO6 unit, now AlO4FPy, has considerably larger quadrupolar parameters [30]. The closeness to the Al(OiPr)3 structure may also be seen by comparing the two 1H-13C CP MAS NMR spectra (Figure 1.10).

14

Functionalized Inorganic Fluorides –160 ppm

100

0

–100

(ppm)

0

–20 –40 –60 –80 –100–120 –140 –160 –180 –200 –220 –240 –260 –280 –300 –320

(ppm)

Figure 1.9 27Al and 19F MAS NMR spectra of Al3(OiPr)8FPy; ( rot ¼ 25 kHz, B0 ¼ 9.4 T) (Reprinted with permission from [30] Copyright (2008) Humboldt University of Berlin.)

80

70

60

50

40

30

20

10

40

30

20

10

(ppm)

70

60

50 (ppm)

Figure 1.10 1H-13C CP MAS spectrum of Al3(OiPr)8FPy (bottom) in comparison with analogous spectrum of its precursor Al(OiPr)3 (top) ( rot ¼ 10kHz, B0 ¼ 9.4 T) (Reprinted with permission from [16] Copyright (2006) Humboldt University, Reprinted with permission from [30] Copyright (2008) Humboldt University of Berlin.)

Those CHO and CH3 groups located on bridged Al-O-Al positions (s. 1, 3 in Figure 1.5), i.e. at the positions > 65 ppm and < 25 ppm, respectively, are especially strongly affected by the new substituents on the central aluminium site. In contrast, the terminal groups are, as expected, less influenced.

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

15

Likewise, the polymeric nature of magnesium methylate results in the bulky Mg6F2(OMe)10(MeOH)14 (Figure 1.11), the synthesis of which can easily be reproduced. In this case the structure is also very similar to that of the fluorine-free compound Mg(OCH3)2 [33]. To preserve the cube-like shape, two fluorine ligands have to be introduced. The high symmetry of the crystal, however, results in one narrow 19F signal at 174 ppm, as shown in [32, Figure 5 therein].

012

011

010 07 08

04 06

Mg3 09

Mg2 03

05

F 01

Mg1

02

Figure 1.11 Crystal structure of Mg6F2(OMe)10(MeOH)14 (Reprinted with permission from [31] Copyright (2008) Wiley-VCH.)

1.2.2.3

Changes in Dry Gels with Progressive Fluorination

Isolated single crystals and their structures only give an idea of very early steps of fluorination, producible only with an understoichiometric fluorine supply. Changes in the local aluminium and fluorine coordination in solutions, sols and thin gels are

16

Functionalized Inorganic Fluorides

highlighted above (Figure 1.7). Drying of such thin and thick wet gels with different Al:F molar ratios leads exclusively to X-ray amorphous materials. However, based on the relevant liquid state NMR spectra (Figure 1.7), fundamental modifications can also be expected for the local coordinations in these solids. Together with the data already presented, a consistent mechanism of the fluorination process is then deducible. Taking all results into account, three stages of the sol-gel fluorination process can be identified, which are shown in the following. The first stage is primarily represented by samples with molar ratios Al to F higher than 1 (meaning low F-contents). Samples with molar ratios 1:1 and 2:3 (stage 2) mark for the solids the changeover to the third stage, the aluminium isopropoxide fluoride xerogels, samples with Al:F ratios lower than 0.5. 1 H-13C CP MAS, 27Al and 19F MAS spectra of dried gels with varying composition, as depicted in Figures 1.12 and 1.13, illustrate these findings. Beside a general shift of the 13C

Al O

Al(OiPr)3

Al

O

Al

O

CH

CH

Al CH3

Al : F

4:1

2:1

1:1

2:3

1:2

1:3 80

70

H

F O

CH

Al

60

50

40 δ13C /ppm

30

H

H

H O CH

20

10 F O

O CH3

Al CH3

Figure 1.12 1H-13C CP MAS NMR spectra (central transitions) of Al(OiPr)3 and aluminium isopropoxide fluoride solids prepared with different molar ratios Al: F as given in the figure ( rot ¼ 10 kHz, B0 ¼ 9.4 T) (Reprinted with permission from [34] Copyright (2009) American Chemical Society.)

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

17

signals to higher fields (CH3 groups) and lower fields (CHO groups), a considerable broadening effect is observable for all bridging CHO – and CH3 groups, and after that, with a higher proportion of fluorine, also a broadening and finally disappearing of the signals of the terminal groups (Figure 1.12) [34]. The spectral changes (peak maxima, line forms) deducible only from the 1H MAS NMR spectra are only minor and the superimposition of the several signals makes it difficult to discuss them separately [34]. 27Al and 19 F MAS NMR spectra however address local changes in the solids during continued fluorination very clearly. As mentioned above, the greatest change can be observed by introducing more and more fluorine and passing the molar ratio Al to F equal to one (see also Figure 1.13). Whereas the spectral features of the initial Al(OiPr)3 are still present in the 27Al MAS NMR spectra of samples with low fluorine content, a rising signal at 38 ppm is detected, which is provoked by a tetrahedrally coordinated aluminium site in the proximity to fluorine as evidenced by 19 27 F Al CP MAS NMR experiments [34]. Their 19F MAS NMR spectra (Figure 1.13) are dominated by very sharp signals (FWHM less than 1 kHz), which indicate ordered (‘crystal-like’) local structures. Most of the species in these solids are more-or-less

Al(OiPr)3

4:1

2:1 –162 –156

–171

–148 –138

–182

1:1

2:3

1:2

1:3 100

75

50

25

0 –25 δ27Al /ppm

–50

–75 –100 –50 –75 –100 –125 –150 –175 –200 –225 –250 δ19F /ppm

Figure 1.13 27Al and 19F MAS NMR spectra of Al(OiPr)3 and aluminium isopropoxide fluoride solids prepared with different molar ratios Al: F as given in the figure ( rot ¼ 25 kHz, B0 ¼ 9.4 T) (Reprinted with permission from [34] Copyright (2009) American Chemical Society.)

18

Functionalized Inorganic Fluorides

isolated; no proximity of the certain F-sites to each other can be stated from 19F-19F spin exchange experiments [34]. The 3QMAS NMR spectra of samples d and e (Figure 1.13) indicate the existence of a set of different AlFx(OiPr)4x – AlFx(OiPr)5x and AlFx(OiPr)6x – species (for the latter x ¼ 3–5) [34]. They are also responsible for the remarkable line-broadening effects in the corresponding fluorine spectra. The existence of certain fourfold and fivefold coordinated AlFx(OiPr)CN-x species as intermediate structures in aluminium isopropoxide fluorides was also unambiguously shown utilizing, for the first time, ultra high-field MAS NMR at magnetic fields B0 up to 21.1 T [35]. Moving to the third stage, a more and more stable network is formed with Al:F ratios equal to 1:2 and 1:3. The amount and spread of fourfold and fivefold coordinated Al-species decreases, ending up with sixfold AlFx(OiPr)6–x species (x ¼ 4 and 5) as deduced from the chemical shift correlation graphs [2628, 36, 37]. The comparison of the development of the intensities of single species with rising fluorination degree with the general development of the certain contributions of the appropriate 19F MAS NMR spectra allows a simple correlation of Al and corresponding F-species. Besides, a variety of possibly terminal fluorine-sites are evident for the highly disordered and amorphous aluminium isopropoxide fluorides in the up-field part of the spectra [35]. 1.2.2.4

Structure of Wet and Dry Aluminium Alkoxide Fluoride Gels AlF2.3OiPr0.7.xiPrOH – A Comparison

An attribution of local structures in wet gelatinous and air-sensitive fluoride gels implies many difficulties, which made the development and testing of inserts for MAS experiments necessary. Alternatively, low-temperature MAS NMR experiments at temperatures below the melting point of the solvent used allow the gel to be filled directly into the rotor [38]. 27Al and 19F MAS NMR spectra of a wet gel recorded at a spinning speed of 10 kHz both in a quartz insert and at low temperature are shown in Figure 1.14 together with the corresponding static spectra. A broad signal around 0 ppm was obtained for the 27Al static NMR spectrum (Figure 1.14a), completely covering the region for AlO4, AlO5, AlO6, AlOxFy or AlF6 species. The very broad static 19F signal is superimposed by narrow lines at –150 ppm and –171 ppm. The latter can be assigned to ‘free’ and mobile F- ions. Rotation at 10 kHz discloses a substantial line narrowing with a 27Al line at –16 ppm with a shoulder and an asymmetric decay in the high-field part. The 19F MAS NMR spectra of the wet gel observed using the quartz insert and applying low temperature are also comparable (Figure 1.14). A broad main signal is visible at –163 ppm together with spinning side bands superimposed by narrow peaks as contributions from the solvent. Comparing the MAS spectra of wet and dry gels, it becomes obvious that most of the structural features are already preformed in the jelly-like gel and in principle conserved and strengthened in the dry gel. This holds for all studied nuclei (19F, 1H, 13C, 27Al [38]). For 27Al this comparison is given in Figure 1.15 [38]. The shoulder and the position of the 27Al central lines as well as the wide spread of spinning side bands are characteristic features for both wet and dry gels. Sixfold coordinated Al-species (AlFxO6-x) are the dominating units in their structure. Second-order quadrupolar broadening is the main line-width factor of the Al signals. Four different structural units could be assigned by additional MQMAS experiments (Figure 1.16, [40]).

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

19

c

27Al

19F

a

–150 ppm –171 ppm

b

static, rotor

static, glass –16 ppm

*

*

*

10 kHz, quartz

* *

10 kHz, quartz * ** * * *

10 kHz, 150 K 1000 800 600

400

* *

10 kHz, 155 K * 200

0

–200 –400 –600 –800 –1000 100 50

0

–50 –100 –150 –200 –250 –300 –350 –400

(ppm)

(ppm)

Figure 1.14 MAS NMR spectra (B0 ¼ 9.4 T) of the wet gel using glass and quartz as insert materials in comparison with the respective static spectra (na: number of accumulations): a) 27Al: rot ¼ 10 kHz, quartz insert (na: 150000); rot ¼ 10 kHz, frozen wet gel in the rotor at 150 K (na: 1000); (b) 19F: rot ¼ 10 kHz, quartz insert (na: 192); frozen wet gel in the rotor at 155 K (na: 48);*: spinning side bands; (c) picture of a wet gel after rotation at 10 kHz in a quartz insert. (Reprinted with permission from [38] Copyright (2007) Elsevier Ltd.)

The reconstruction of the 27Al MAS NMR spectrum of the xerogel was possible with the four sites (A-D, Figure 1.16) obtained by analysis of the 3QMAS spectrum. [40]. These sites are attributed to AlF0-2O6-4 units (A), AlF4-5O2-1 units (38 %, B and C) and AlF5O1 units (53 % D) in the network [27, 40]. The existence of already immobilized –OiPr-groups, incorporated into or associated on the network, can be unambiguously proven for the wet gel by 13C CP MAS experiments [38]. In addition, the more rigid structure of the dry gel allows distinctions to be made between matrix groups and associated immobilized iPrOH molecules [38, 41]. The main structural features of the dried gel are already built up in the jelly-like gel and change only little in aging and drying processes. 1.2.2.5

Network Formation with Increasing Fluorine Content

Bearing in mind all NMR spectra recorded for sols, thin wet gels and dried gels (Figures 1.7, 1.12–1.16) in comparison with those of solid and liquid educts, the following mechanism can be derived for the fluorolytic sol-gel reaction. In the first step of the synthesis route of HS-AlF3, which is usually applied, the tetrameric structure of Al(OiPr)3

20

Functionalized Inorganic Fluorides

wet gel 160 K

a

100

75

50

25

0

–25 –50 (ppm)

–75 –100 –125 –150

Xerogel

b

νrot = 25 kHz

100

75

50

25

0

–25

–50

–75 –100 –125 –150

(ppm)

Figure 1.15 27Al MAS NMR spectra of the wet and dried gel (xerogel) with focus on the central region (400 MHz spectrometer; na: number of accumulations): (a) comparison of 27Al MAS NMR spectra of the frozen wet gel with (solid line, 160 K) and without (dotted line, 155 K) 19F (cw) decoupling at cryo-temperatures, for both: rot ¼ 10 kHz, na: 1000; (b) possible different contributions to the spectrum of the xerogel including the n ¼ 0 spinning sideband of the satellite transition: (- - -). (s. also Neuville, Massiot [39]); experimental spectrum (——); simulated spectrum () and decomposition ( - - -) (see also Table 2 in [38]). (Reprinted with permission from [38] Copyright (2007) Elsevier Ltd.)

as dominating species (Figure 1.5, 1) represents the starting point for the further reaction pathway and finally for the formation of the gel network. The fluorination begins with a protonation of bridging isopropoxide groups, which makes a subsequent attack of fluorine ions or HF to central Al atoms easy. As result, a substitution of protonated isopropoxide groups by fluorine occurs. At early fluorination states (state 1, Scheme 1.4) unconverted Al(OiPr)3-molecules exist in the sols and gels to some degree – either in its tetrameric or in its linear trimeric form. In the following the latter may be partly fluorinated and, if stabilizing donors (D) are

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

δMQMAS(27Al) in ppm

–40

qc

cs

D –20

21

C

A

0 20 B 40 20

0

–20 –40 δMAS(27Al) in ppm

–60

Figure 1.16 3QMAS spectrum of the aluminium alkoxide fluoride gel (Al:F ¼ 1:3) acquired at 14.1 T (Reprinted with permission from [40] Copyright (2009) American Chemical Society.)

accessible, even crystals of Al3(OiPr)8F•D can be isolated (Scheme 1.4, 3). Nevertheless, for aluminium isopropoxide fluorides resulting from sols without donor molecules the formation of ordered similar species Al3(OiPr)8F•iPrOH is presumable as evidenced, for example, by the corresponding sharp signals centred at –160 ppm and –171 ppm (19F). The early formation of tetrahedrally coordinated species AlF4 can be demonstrated for solids prepared with low fluorine supply. This species does not seem to exist in the corresponding sols. Instead, it is plausible that these species are stabilized as solvated species AlFx(HOiPr)6x3x (x ¼ 4) in the sols. Drying results in a partial loss of the solvent molecules and the formation of ‘incorporated’ AlF4-species. Additionally, further linking processes of the AlF4(HOiPr)2 species lead to the formation of bigger units (beginning of a gelous network), which, as a consequence, are predominantly built of AlFx((HOiPr)6-xoctahedra (x ¼ 3 to 5). The rise of the F-content (molar ratio Al: F < 1) leads to an irregularly strongly distorted and disordered solid. The corresponding sol consists of a loose network, imaginable in solution as presented earlier in [23]. Vacuum drying leads to cleavage of iPrOH molecules coordinating to Al-species existing in these sols and gels, forming a variety of fourfold, fivefold and sixfold coordinated Al-species like AlFx(H)OiPr)4x, AlFx((H)OiPr)5x and AlFx((H)OiPr)6x as identified by 3QMAS and ultra high-field MAS NMR measurements [34, 35]. The units are sterically separated; as in the 19F-19F EXSY NMR experiments, nearly no spin exchange could be observed ([34], state 2, s. Scheme 1.5). Finally this gelous network is strengthened and stabilized by cross-linking with higher fluorine supply. Vacuum drying leads then to xerogels that consist predominantly of sixfold coordinated AlFx((H)OiPr)6x species (x: 35)[38, 40]. Coordinating solvent molecules are stabilized by the formation of H-bridges. The local structures of the xerogel are preformed in the wet gel. Nevertheless, ordered local structures are also observable for the xerogels. The derived mechanism is schematically depicted in Schemes 1.4 and 1.5.

22

Functionalized Inorganic Fluorides O

State 1

O side reaction

O

Al

O Al O

O

if donor present isolable

O Al O F D

Al O

3

species present in isopropoxide fluoride sols with a ratio F: Al lower 1… O

Al O O

O O

O

O O Al O O O Al O O 1

Al O

Al O O O O Al O O

O*

F

Al

F

F F

*O

O*

*O Al

2

F F

F

F F

F

O Al

Al F

F

F

F

– O*

F O

O F O Al F * F F

F Al

F

F

…and in the corresponding solids.

The loss of solvent leads to formation of AlF4 species along with un converted Al(OiPr)3 1 and cross-linked AlFx(OiPr)6–x species.

F

t

F O* Al F O F F Al O O Al O* F F F * F F

F

Al

F

F

Scheme 1.4 Derived mechanism for the beginning sol-gel fluorination process with low fluorine supply (O*: iPrOH; O: -O-iPr). (Reproduced from [34] by permission of the American Chemical Society.) State 2

wet sol:

F

F *O

O

F

Al

solid:

F

F

F F F O F Al *O Al O F F F * Al F F F F F F Al F O F F O* Al F O F F Al O O Al O* F F F * F F

and further

– O*

O Al

O

F F

F

Al

F F F F F F F O F Al Al F *O Al F O F *O Al F F * Al F F F F F F F F F Al F O F F Al F F O O* Al F F F Al Al O F O F O* Al F F * F F F F Al O O Al F F F * F F *O

irregular solid

O

minor ? O Al O

Al O O O O Al O O* F 3` (D = O*)

indications for formationof a similar compound like 3

Scheme 1.5 Derived mechanism for a progressive fluorination degree (reproduced from [34] by permission of the American Chemical Society)

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

1.2.2.6

23

Structural Characterization of Sol-Gel Derived High-surface Fluorides

1.2.2.6.1 Vibrational Spectroscopy of HS-AlF3 The IR transmission spectra of HS-AlF3 were recorded and compared with those of aluminium isopropylate and of dry aluminium fluoride gel – i.e. the Al moieties involved in HS-AlF3 synthesis (Figure 1.17) [8], as well as of crystalline AlF3 phases (Figures 1.18 and 1.19) [42]. In Figure 1.17 the CH3 and CH as well OH absorption bands of isopropoxide and/or isopropanol can be seen at wave numbers higher than 2800 cm-1 in the spectra of aluminium isopropylate and of dry aluminium fluoride gel, whereas the broad band in the HS-AlF3 spectrum can be attributed (i) to water adsorbed in course of sample preparation due to the high Lewis acidity, and (ii) to water residues present in the bulk as a consequence of the formation of diisopropylether at postfluorination. At low wave numbers there are the absorption bands of Al-O and Al-F valence vibrations (360–700 cm-1) and of C-C frame vibrations of the organic constituents. These spectra show the disappearance of organic components converting the dry gel, the ‘precursor’, into HS-AlF3.

Transmission

HS-AlF3

precursor

Al(Oi Pr)3

4000

3500

3000

2500

2000

Wave number

1500

1000

500

(cm–1)

Figure 1.17 IR spectra of Al(OiPr)3, the AlF2.3OiPr0.7.xiPrOH gel, and HS-AlF3 (Reprinted with permission from [8] Copyright (2007) Springer Science þ Business Media.)

The amorphous, highly disordered state of HS-AlF3 follows convincingly from a comparison of its IR spectra with those of -AlF3 shown in Figure 1.18. Since crystalline -AlF3 also gives a well resolved Raman spectrum (Figure 1.19) all the distorted X-ray amorphous phases of HS-AlF3 do not yield any useful information. The very broad peaks in the IR-spectrum of HS-AlF3 are obviously the consequence of its amorphous state and can be interpreted as superposition of many unresolved peaks, also covering the range of vibration bands of the - and -AlF3 phases. 1.2.2.6.2 X-ray Diffraction and TEM Investigations of HS-AlF3 In contrast to vibrational spectroscopy, the X-ray diffraction patterns of HS-AlF3 and of its precursor, the AlF2.3OiPr0.7.xiPrOH dry gel, show no peak and therefore only indicate the

24

Functionalized Inorganic Fluorides

HS-AlF3

Transmission

β-AlF3

α-AlF3

0

200

400 600 800 Wavenumber/cm–1

1000

1200

Figure 1.18 IR-transmission spectra of a-AlF3, -AlF3, and HS-AlF3 (Reprinted with permission from [42] Copyright (2007) American Chemical Society.) 157

30 000

20 000

10 000 5000

478

383

15 000

96

Intensity [arb.units]

25 000

100

x25

200

300

400 500 Wavenumber [cm–1]

600

700

800

Figure 1.19 Raman spectrum of a-AlF3 (rhombohedral phase) (Reprinted with permission from [42] Copyright (2007) American Chemical Society.)

X-ray amorphous state of these samples. Only after heating above the crystallization temperature of about 836 K do peaks corresponding to -AlF3 become apparent (Figure 1.20). In the TEM micrograph (Figure 1.21) both the agglomeration of nano-particles to microparticles as well as the partial crystalinity of the nano-particles can be seen. 1.2.2.6.3 Solid State NMR of HS-AlF3 Subsequent drying of the xerogel in vacuum results in a noticeable loss of –OiPr-groups and associated iPrOH molecules, which can easily be followed by 1H-13C CP MAS

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

25

Intensity (au)

(a)-precursor (b)-precursor in N2 at 350 °C (c)-HS-AlF3 (d)-precursor in N2 at 700 °C (α-AlF3)

(d) (c) (b) (a)

10

20

30

40

50

60

2θ

Figure 1.20 X-ray powder difractograms of (a) AlF2.3OiPr0.7.xiPrOH dry gel, (b) AlF2.3OiPr0.7.xiPrOH dry gel heated to 623K, (c) HS-AlF3, and (d) AlF2.3OiPr0.7.xiPrOH dry gel heated to 973 K (Reprinted with permission from [8] Copyright (2007) Springer Science þ Business Media.)

Figure 1.21 TEM micrograph of HS-AlF3 (Reprinted with permission from [6] Copyright (2003) Wiley-VCH.)

experiments [44]. Both 19F and 27Al MAS NMR spectra exhibit after removal of the organic components the typical shape as recorded for HS-AlF3. The maximum broad 19F MAS NMR signal lies typically at –165 ppm [41]. The effect of such a drying process is shown exemplarily for 27Al in Figure 1.22. The spectrum on the bottom (Figure 1.22) is almost identical with the central line usually measured for HS-AlF3. In contrast to the xerogel (AlF2.3OiPr0.7.xiPrOH) the 27Al MAS NMR signal line widths do not narrow

26

Functionalized Inorganic Fluorides

much upon increase of the applied magnetic field. Therefore the comparison of MQMAS experiments at different magnetic field strength is most important for an assignment of various species. Although the data analysis is very difficult, the presence of various signals was revealed [40]. A 3QMAS spectrum of HS-AlF3 is given in Figure 1.23.

70 °C, 2 h

100 °C, 2 h -–15 ppm

150 °C, 2 h

170 °C, 4 h 100

50

0

–50

–100

–150

(ppm)

Figure 1.22 27Al MAS NMR spectra (central lines) of a AlF2.3OiPr0.7.xiPrOH gel thermally dried in a vacuum at different temperatures ( rot ¼ 25 kHz, B0 ¼ 9,4 T) (Reprinted with permission from [44] Copyright (2009) Humboldt University of Berlin.) –40 qc δMQMAS(27Al) in ppm

B'

cs

–20 A' 0 20 40 20

0

–40 –20 δMAS(27Al) in ppm

–60

Figure 1.23 3QMAS NMR spectrum of HS-AlF3 acquired at 19.9 T (Reprinted with permission from [40] Copyright (2009) American Chemical Society.)

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

27

The line shape analysis of the 27Al MAS NMR spectrum of HS-AlF3 required more than the two Al-sites extracted from the 3QMAS spectrum (Figure 1.23). A successful fit was only possible by implementation of two further Al-sites. For all of them the simulation required the Czjzek-model, i.e. distributions of quadrupolar parameters were used [40]. Now 65 % of the integral intensity of the Al signal belongs to AlF5-6O1-0 and AlF6 units, respectively. Residues of AlF4-5O2-1 units (25 %) and AlF0-2O6-4 units (11 %) are still present [40].

1.2.3

Exploring Properties

The most remarkable property of HS-AlF3 is its outstanding Lewis acidity [6, 8, 9, 45], which is far higher than that of AlCl3, as will be shown later. It is interesting to note that on a theoretical basis for an isolated ‘AlF3-molecule’ already a Lewis acidity was predicted ranking among the highest ones at all [46]. In the following a variety of investigations showing the very high Lewis acidity of HS-AlF3 will be presented and discussed.

1.2.3.1

Adsorption of Probe Molecules

The strength and/or nature and/or amount of acid sites accessible at the surface of a solid can principally be determined measuring the interaction with a basic probe molecule. The higher the acidity of the solid the lower can be the basicity of the probe and vice versa. For investigations of HS-AlF3, pyridine and its derivatives, ammonia and carbon monoxide, were employed. Their interactions with the solid have been analysed studying in case of ammonia desorption with increasing temperature, Temperature Programmed Desorption TPD, and in case of the other probes via IR spectroscopy. NH3-TPD: HS-AlF3 was saturated with gaseous NH3 followed by flushing with N2. Upon gradually heating the sample up to 773 K, i.e. below the crystallization temperature, the adsorbed NH3 is gradually released as shown in Figure 1.24. Compared to the well-known solid Lewis acid -AlF3 the desorption from HS-AlF3 occurs up to much higher temperatures corresponding to a much higher Lewis acidity, and the total amount of NH3 is also much higher, indicating a higher number of acidic sites per gram. Pyridine adsorption: The chemical nature of the acid sites can be seen from photoacoustic infrared spectra (PAS) of adsorbed pyridine (Figure 1.25). Frequencies and intensities of the IR bands show for HS-AlF3 (almost) only Lewis acid sites, whereas a HS-AlF3 sample showed after treatment with HF predominantly Brønsted acid sites (Figure 1.25b). Carbon monoxide adsorption: Carbon monoxide behaves towards a strong acid as a weak base the interaction of which can be investigated monitoring the CO stretching region of the IR absorption spectrum of the absorbed CO. The stronger the acid the more is the CO IR frequency blue-shifted. HS-AlF3 shows the strongest blue shift ever reported for a solid acid (for details see Chapter 3) indicating that it is, next to ACF (aluminium chlorofluoride), the strongest solid acid of all [47].

28

Functionalized Inorganic Fluorides HS-AIF3 beta-AIF3

35 000 30 000 25 000 20 000 15 000 10 000 5000 0 –5000 0

100

200

300

400

500

Temperature/°C

Figure 1.24 NH3-TPD of HS-AlF3 in comparison to -AlF3 (Reprinted with permission from [7] Copyright (2007) Royal Society of Chemistry.)

Figure 1.25 PAS of pyridine adsorbed on (a) HS-AlF3, (b) HS-AlF3 exposed to HF, (c) HS-AlF3 prepared by post-fluorination with HF, and (d)sample as in(c) additionally heated in CHClF2/N2 flow; peaks at (B) are indicative for Brønsted acidity, and peaks at (L) for Lewis acidity (Reprinted with permission from [8], Figure, Copyright (2007) Springer Science þ Business Media.)

1.2.3.2

Catalytic Test Reactions

Reactions which have to be catalysed by a Lewis acid to proceed can be used as testreaction for assessment of the acidity of a material under study. Equations (1.5) to (1.8) show four reactions, the use of which has been reported [8, 45]: 5 CCl2 F2 ! CCl3 F þ 3 CCl3 F þ CCl4

(1:5)

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

29

3 CHClF2 ! CHCl3 þ 2 CHF3

(1:6)

CBrF2 CBrFCF3 ! CF3 CBr2 CF3

(1:7)

CCl2 FCClF2 ! CCl3 CF3

(1:8)

The educts of the reactions in Equations (1.5) and (1.6) were used as postfluorination agents in the course of high surface metal fluoride preparation, especially that of HS-AlF3. As soon as first parts of the dry Al-F-gel are converted into HS-AlF3, a dismutation reaction according to Equations (1.5) or (1.6) starts, which can be detected by GC. Thus, these fluorinating agents act conveniently as detectors for Lewis acidity. However, the reactions in Equations (1.5) and (1.6) already proceed in the presence of comparable weak Lewis acids; they are therefore no measure of high acidity. The isomerization of CBrF2CBrFCF3 (Equation 1.7), on the other hand, proceeds only under the catalytic action of the strongest known Lewis acids, i.e. antimony pentafluoride or aluminium chlorofluoride, and is also catalysed by HS-AlF3 but not by AlCl3 [8, 45]. Thus, the isomerization reaction gives convincing evidence for the exceptional high Lewis acidity of HS-AlF3. As AlCl3 is widely used in organic synthesis as a Lewis acidic catalyst a further comparison of its Lewis acidity with that of HS-AlF3 was of interest. Studying the isomerization of CCl2FCClF2 (Equation 1.8) as test reaction, AlCl3, HS-AlF3, ACF, -AlF3 and -AlF3 have been compared regarding their respective catalytic activity. It was found that ACF and HS-AlF3 were catalytically very active, and -AlF3 and -AlF3 not at all, as expected. Surprisingly, AlCl3 was also not primarily active but became active only under conditions and after some time as was needed to convert AlCl3 into ACF. The experimental result was interpreted based on theoretical investigations assuming that under-coordinated Al atoms, which are a result of the high degree of disorder, are responsible for the Lewis acidity of HS-AlF3 [45]. Radiotracer investigations: Radiotracer experiments, which also gave evidence for the exceptional Lewis acidity of HS-AlF3, are discussed in Chapter 3.

1.2.4 1.2.4.1

Possible Fields of Application Range of Metal Fluorides Obtainable via Sol-Gel Fluorination

The fluorolytic sol-gel synthesis of metal fluorides was originally developed and explored for aluminium fluoride, which was a piece of luck since both the stepwise synthesis and the properties of HS-AlF3 showed the influence of the new synthesis process. Thus, almost immediately after exploration of HS-AlF3 other binary metal fluorides have been similarly prepared and attempted syntheses of more complex systems started. Binary metal fluorides: Many binary metal fluorides have been prepared via sol-gel fluorolysis [4, 7]. The applicability of the synthesis process is primarily limited by the ready availability of soluble metal alkoxides. However, the synthesis of metal fluorides, of which the metal ions are very weak Lewis acids, typically does not result in sol formation but finely dispersed solid fluorides, the XRD of which reflect the aimed-for compounds.

30

Functionalized Inorganic Fluorides

Thus, upon fluorolysis of the tert-butoxides of Li, Na, K and Cs in THF solution, only with LiOtBu and NaOtBu did gel formation occur whereas with K- and Cs- alkoxide immediate precipitation was observed. The dried products of all these alkali metal ions reflected the corresponding fluorides in XRD and also gave some indications of the respective hydrogenfluorides due to their higher thermodynamic stability [48]. The most thoroughly investigated example besides HS-AlF3 is HS-MgF2 [11]. Other high surface area alkali earth fluorides prepared are HS-CaF2 and HS-BaF2 [49]. Mixed metal fluorides: Prompted by a hypothetical model by Tanabe explaining the Lewis acidity of guest-host metal oxide systems [50] and its adaptation for fluoride systems [51], guest-host mixed metal fluoride systems with HS-MgF2 as host have been prepared (see also Chapter 3). If the radii are comparable, the metal ions guest ions will probably occupy places of Mg2þ ions. If the guest ion has a higher positive charge than Mg2þ, Lewis acidity should be created in accordance with the model. This way, solid potential metal fluoride catalysts with tunable Lewis acidity are accessible. Thus, with up to 20 mol % of Fe3þ, Ga3þ, V3þ, In3þ, and Cr3þ as guest components, solid solutions with HS-MgF2 as host have been prepared [52–54]. The synthesis of such systems basically followed the fluorolytic sol-gel route described above, however, in some experiments compounds other than alkoxides have been employed as guest components to be added to the Mg alkoxide solution. Analyses of these mixed systems by XRD, 19F MAS NMR and photoelectron spectroscopy showed no evidence of the guest compounds but gave only of the typical spectral data of MgF2, which however, were not identical with those of crystalline MgF2. Obviously the guest metal ions were incorporated into the MgF2 lattice as expected. ESCA investigations in the Cr3þ/MgF2 system (Figure 1.26) showed, for both Mg 1s and F 1s, binding energies near those of pure MgF2 [54]. A shift that is only small to lower binding energies gives evidence that the chemical environment of Mg and F in the mixed systems is influenced by Cr3þ. Complex metal fluorides: Complex metal fluorides are especially interesting compounds because of their physical properties, which are suitable for applications such as lasers. Conventional syntheses typically employ thermal methods starting from the respective metal fluorides. Complex metal fluorides can also be conveniently prepared via sol-gel fluorination. The sol-gel syntheses start similar to that of host-guest systems from mixtures of the respective metal alkoxides but in the respective stoichiometric ratio, which are subjected to fluorolysis. Contrary to the host-guest systems, where the host lattice is preserved, complex systems, i.e. fluorometallates, have their specific crystal structure different from those of the respective binary metal fluorides they are composed of. Examples of complex metal fluorides prepared via sol-gel fluorolysis are Li3AlF6, Na3AlF6, K3AlF6, KAlF4, CsAlF4, LiNa2AlF6 [55] and BaAlF5, K2MgF4 and BaMgF4 [49]. Since Al in AlF6 units is ideally shielded, such compounds do not show any Lewis acidity and are chemically very stable making them suited especially for nonchemical applications such as protective coating (see below). 1.2.4.2

Application Consequences of the Sol-Gel Synthesis

Specific chemical and physical properties of intermediate states of the sol-gel fluorolysis can be utilized for quite different modifications and applications. Upon reaction of metal

31

655

1339

1179

2482

654

1338

1178

2481

653

1337 MgF2 CrClMg15 CrOMg8 CrOMg15 CrOMg25 CrOMg50 crAcMg15 CrAcMg40

652 651 650 689

688

687 686 F ls (eV)

685

1336 1335 1334 684

Mg KLL (eV)

F KLL (eV)

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

1177

2480 MgF2 CrClMg15 CrOMg8 CrOMg15 CrOMg25 CrOMg50 crAcMg15 CrAcMg40

1176 1175 1174 1308

1307

1305 1306 Mg ls (eV)

1304

2479 2478

2477 1303

Figure 1.26 Chemical state plots for F 1s and Mg 1s orbitals of chromium(III)-doped MgF2 prepared from CrO3 (CrO), CrCl3 (CrCl) or Cr-acetat (CrAc) in different concentrations (Reprinted with permission from [54] Copyright (2005) Elsevier Ltd.)

alkoxides with understoichiometric amounts of hydrogen fluoride a defined part of the alkoxide groups remain and may be used for chemical modifications of the high surface metal fluorides. Such modifications include partial substitution of F by OH, chemical immobilization of metal oxide oxidation catalyst onto the metal fluoride and binding organic groups to the inorganic metal fluoride. Introduction of highly dispersed noble metals on metal fluoride is also possible. The sol state, i.e. the colloidal solution of a metal fluoride in a nonaqueous solvent, opens up the possibility of many interesting applications. Most important is the ability to prepare different coatings for very different applications. Thus, catalytic active metal fluorides can be conveniently deposited on supports, or thin metal fluoride layers, can be easily prepared for, e.g., optical or mechanical applications. Nano-sized metal fluorides: An optically clear metal fluoride sol contains particles the diameter of which is in the range of the wavelength of visible light. Therefore, nano-sized metal fluoride particles can be obtained from such sols, which are normally agglomerated, after removal of the solvent at moderate temperatures. The size of the individual particles has been shown by TEM and was derived from their XRD pattern to be below 5 nm [6, 49]. Such tiny particles represent a higher state of energy compared to more bulky material. As a consequence these materials can be easily pressed at room temperature to dense glasses. Thus, transparent glasses have been obtained from MgF2, CaF2, BaMgF4 [49] and Rb2NaAlF6 [56] employing pressures up to 1.1 GPa at room temperature. Nano-sized metal fluorides have a promising perspective as inorganic component of organic polymers. Due to their small size the particles are invisible even in transparent polymers but may, in possibly decisive ways, change physical properties such as dielectric constant, index of refraction or mechanical properties. However, preventing agglomeration of these nano particles is a challenging task, but one that is not easy to accomplish and that needs to be further developed. Oxide fluorides: The amount of alkoxide groups remaining after partial fluorination can be tuned over a wide range. These OR groups can subsequently be hydrolysed, i.e. substituted by OH, or thermally split off, also resulting in OH and, due to condensation reaction, in the formation of oxide groups homogeneously distributed within the fluoride matrix. The metal fluorides modified this way exhibit not only Lewis acidity but also some Brønsted acidity or

32

Functionalized Inorganic Fluorides

even basicity. Such bifunctional materials can be valuable solid catalysts (see below). This synthesis principle has been realized for magnesium oxide/hydroxide fluorides within a broad range of compositions [57]. It was found that preparations with low fluorine content, with nominal composition Mg(OH)1.2F0.8 and Mg(OH)0.8F1.2, were almost X-ray amorphous even after calcinations at 623 K whereas with higher F content the patterns of MgF2 appeared. In 19F MAS NMR there was a similar trend to be seen, the low F materials gave broad, complex signals, indicating the presence of many different fluorine species, which became more and more sharp and like those in MgF2 with increasing F content. The higher electronegativity of F compared to OH resulted with increasing F content in a reduced electron density at Mg indicated in increasing Mg 1s binding energy as seen in XPS analysis. Magnesium fluoride-based bifunctional materials have been successfully employed as heterogeneous catalysts for quite different reactions (see below) [57–59]. Metal oxides linked to metal fluorides: The performance of a solid catalyst depends to a substantial degree on its surface area, which can often be increased by suitable support, and then also on the chemical and texture properties of the support. Vanadium oxide-based catalysts are very useful in oxidation reactions because of the easy change between different oxidation states. In case of the technically important partial oxidation of organic compounds, a limited oxidation activity of the catalyst is needed to prevent total oxidation to CO / CO2 and H2O. As a consequence the organic educt has to be activated by, e.g., protonation due to Brønsted acidity of a catalyst support. By using a Lewis acidic support a more gentle activation of the educts should be possible resulting in less or no total oxidation. The sol-gel synthesis provides exquisite conditions for a very evenly distributed and highly dispersed VOx based catalyst deposited on Lewis acidic metal fluoride supports. Employing the principle of under-stoichiometric fluorination there will be OR groups or, after hydrolysis, OH groups available for chemical anchoring of the VOx species. This way aluminium fluoride-supported VOx catalysts with outstanding performance have been prepared and thoroughly characterized [60]. Noble_metals supported by high surface area metal fluorides: Noble metals especially platinum and palladium are useful catalysts for many reactions such as, e.g., hydrogenations. They are used in a highly dispersed state deposited on suitable supports. For use in hydrodehalogenation reactions the support has to be stable against attack by the hydrogen halogenide formed in course of the reaction. That requirement can be met using metal fluorides as support. MgF2 and AlF3 were often proposed [61]. For the preparation of very tiny noble metal particles supported on a metal fluoride, which has a large surface, the sol-gel fluorination provides excellent conditions. Starting from the metal alkoxide solution or from the already formed metal fluoride sol an organic solution of a suitable noble metal compound, such as the acetyl acetonate, is added and the mixture is worked up as normal for high surface metal fluoride preparation. Only an additional reduction step is necessary to obtain the catalyst. For a Pd0/CaF2 catalyst prepared accordingly the high Pd dispersion is shown in Figure 1.27. Organically modified metal fluorides: Silicon oxide and also metal oxide based inorganicorganic hybrid materials have received broad interest academically and also for technical applications [62]. The development of such hybrids aims to combine useful properties of the inorganic part, mostly the basis of the hybrid, with those of the organic part, whereby the two parts are chemically bound to one another. Basically, two types of synthesis routes

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

33

Figure 1.27 TEM micrograph of a Pd0/CaF2 catalyst showing Pd particles (black spots) of £10nm diameter (Reprinted with permission from [61] Copyright (2008) Royal Society of Chemistry.)

are used. Both can be adapted to metal fluoride systems. One route starts from organically modified molecular alkoxidic precursors, which are subjected to sol-gel hydrolysis; the other starts from preformed nanoscopic or macroscopic particles to which an organic moiety becomes linked. To transpose these synthesis principles to metal fluoride systems, the sol-gel fluorolysis can be conveniently employed attaching either an organic moiety via a HF stable linkage to metal alkoxides, which are consequently subjected to sol-gel fluorolysis. Another route comprises at first formation of colloidal nanoscopic particles by preparing a sol with under-stoichiometric amounts of HF, to which then organic moieties can be linked. First experiments with phenylphosphonate modified aluminium alkoxide have proved successful [30, 63]. Coatings: The sol state as primary result of the sol-gel fluorination process is very useful, because it consists of nano-sized, colloidal metal fluoride particles homogeneously dispersed in a nonaqueous liquid. Simply by applying the sol on any solid material wettable by the solvent and subsequent drying a metal fluoride layer is on the surface of the material obtained. Depending on the required properties and intended purposes of the layer and on the texture and geometry of the solid material, different methods can be used for the sol application such as, e.g., soaking, dipping, spin coating etc. Catalysis: For technical reasons it can be necessary to have a catalytic active solid material supported to improve, e.g., its mechanical stability and reduce its flow resistance when used in a flow reactor. The sol-gel fluorination synthesis provides a convenient way for depositing high surface area metal fluorides on supports. For example, HS-AlF3, which as fine powder makes problems when used as catalyst in flow systems, could be supported by g-Al2O3 whereby its Lewis acidity and consequently its catalytic activity remains almost unchanged [64]. For other catalytic applications, like micro-reactor techniques, deposition of catalytically active thin layers of metal fluorides is also of interest.

34

Functionalized Inorganic Fluorides

Optics: Probably the most important field of application for thin metal fluoride layers is in physics, especially in optics. A reduction of the reflexion of light at the surface of glass, well known for optical devices such as lenses for glasses and cameras, is even more important for solar energy utilization. Antireflective systems are formed of alternating layers of transparent low and high refractive index materials. However, even with a single layer, a very efficient antireflective system is possible in principle. For a wavelength g and an antireflective layer of thickness g/4 the reflexions at the air/layer surface and at the layer/glass interface have a phase-shift to each other of g/2, that is the precondition for extinction. Total extinction is only possible when the two reflexions are of the same intensity. This can be reached when the index of refraction of the antireflective layer is equal to the geometric mean of air (n ¼ 1.0) and glass (n ¼ 1.5), i.e. for n ¼ 1.225. Typically, the refractive indices of metal oxides are even higher than that of glass, whereas some metal fluorides such as MgF2 (n500 ¼ 1.38), AlF3 (n500 ¼ 1.35) and Na3AlF6 (n500 ¼ 1.33) have distinctly lower indices of refraction, although not as low as 1.225, together with an excellent transparency within a broad range of wavelengths. Consequently, MgF2 thin layers have already found much interest [65]. Methanolic MgF2 sols prepared as described above have been used for spin-coating of planar surfaces [66]. AFM investigations revealed that the layers obtained after drying consist of densely packed particles of 10 to 20 nm diameter as shown in Figure 1.28. The single or multiple MgF2 thin layers showed excellent homogeneity concerning thickness and index of refraction over the experimental range of 5 cm [67].

Figure 1.28 AFM image of a 3-fold deposited MgF2 layer on a silicon wafer, after calcinations at 300 C (area 1 x 1mm2). (Reprinted with permission from [67] Copyright (2008) Wiley-VCH.)

Protective coating: The facileness of preparing metal fluoride layers from nonaqueous metal fluoride sols also reduces the threshold of their use for protective coating. Homogeneous, mechanically stable metal fluoride layers can protect against UV radiation,

Sol-Gel Synthesis of Nano-Scaled Metal Fluorides – Mechanism and Properties

35

chemical and also mechanical impact, and can form a barrier against microbial attack. A CaF2 addition to lithium grease has already proved useful in the lab for friction reduction and extreme pressure properties making CaF2 layers, likewise, interesting. In conclusion, the new access towards nanoscopic metal fluorides via this recently developed fluorolytic sol-gel synthesis route opens a wide range of applications for metal fluorides due to the distinctive different properties of these nano materials.

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[38] R. Ko¨nig, G. Scholz and E. Kemnitz, New inserts and low temperature: two strategies to overcome the bottleneck in MAS NMR on wet gels, Solid State NMR, 32, 78–88 (2007). [39] D. R. Neuville, L. Cormier and D. Massiot, Al environment in tectosilicate and peraluminous glasses: A 27Al MQ-MAS NMR, Raman and XANES investigation, Geochim. et Cosmochim. Acta, 68, 5071–5079 (2004). [40] A. Pawlik, R. Ko¨nig, G. Scholz, E. Kemnitz, G. Brunklaus, M. Bertmer and C. Ja¨ger, Access to local structures of HS-AlF3 and its precursor determined by high resolution solid-state NMR, J. Phys. Chem. C, 113, 16674–16680 (2009). [41] A. Pawlik, R. Ko¨nig, G. Scholz, E. Kemnitz and C. Ja¨ger, HS-AlF3 and its precursor: the structure of amorphous systems investigated by 19F and 1H solid state NMR, in preparation, J. Phys. Chem. C, (2010). [42] U. Groß, St. Ru¨diger, E. Kemnitz, K.-W. Brzezinka, S. Mukhopadhyay, C. L. Bailey, A. Wander and N. M. Harrison, Vibrational analysis study of aluminium trifluoride phases, J. Phys. Chem. A, 111, 5813–5819 (2007). [43] S. Chaudhuri, P. Chupas, B. J. Morgan, P. A. Madden and C. Grey, An atomistic MD simulation and pair-distribution-function study of disorder and reactivity of -AlF3 nanoparticles, Phys. Chem. Chem. Phys., 8, 5045–5055 (2006). [44] R. Ko¨nig, Lokale Strukturen nanoskopischer Aluminiumalkoxidfluoride und chemisch verwandter kristalliner Verbindungen, PhD thesis, Humboldt University of Berlin, (2009). [45] J. K. Murthy, U. Groß, St. Ru¨diger, V. V. Rao, V. V. Kumar, A. Wander, C. L. Bailey, N. M. Harrison and E. Kemnitz, Aluminium chloride as a solid is not a strong Lewis acid, J. Phys. Chem. B, 110, 8314–8319 (2006). [46] K. O. Christe, D. A. Dixon, D. McLemore, W. W. Wilson, J. A. Sheehy and D. A. Boatz, On a quantitative scale for Lewis acidity and recent progress in polynitrogen chemistry, J. Fluorine Chem., 101, 151–153 (2000). [47] Th. Krahl. A. Vimont, G. Eltanany, M. Daturi and E. Kemnitz, Determination of the acidity of high surface AlF3 by IR spectroscopy of adsorbed CO probe molecules, J. Phys. Chem. C, 111, 18317–18325 (2007). [48] M. Ahrens, G. Scholz and E. Kemnitz, Synthesis and crystal structure of RbKLiAlF6 – the first Al-Elpasolite with three different alkali metals, Z. Anorg. Allg. Chem. 634, 2978–2981 (2008). [49] U. Groß, St. Ru¨diger and E. Kemnitz, Alkaline earth fluorides and their complexes: A sol-gel fluorination study, Solid State Sci. 9, 838–842 (2007). [50] K. Tanabe, T. Sumiyoshi, K. Shibata, T. Kiyoura and J. Kitagawa, A new hypothesis regarding the surface acidity of binary metal oxides, Bull. Chem. Soc. Japan, 47, 1064–1066 (1974). [51] E. Kemnitz, Y. Zhu and B. Adamczyk, Enhanced Lewis acidity by aliovalent cation doping in metal fluorides, J. Fluorine Chem., 114, 163–170 (2002). ¨ nveren and E. Kemnitz, Mixed metal fluorides as [52] J. Krishna Murthy, U. Groß, St. Ru¨diger, E. U doped Lewis acidic catalysts systems: a comparative study involving novel high surface area metal fluorides, J. Fluorine Chem., 125, 937–949 (2004). [53] J. Krishna Murthy, U. Groß, St. Ru¨diger and E. Kemnitz, FeF3/MgF2: novel Lewis acidic catalyst systems, Appl. Catal., A, 278, 133–138 (2004). ¨ nveren, W. Unger and E. Kemnitz, Synthesis and [54] J. Krishna Murthy, U. Groß, St. Ru¨diger, E.U characterization of chromium(III)-doped magnesium fluoride catalysts, Appl. Catal., A., 282, 85–91 (2005). [55] M. Ahrens, G. Scholz, M. Feist and E. Kemnitz, Application of an alkoxide sol-gel route for the preparation of complex fluorides of the MAlF4 (M ¼ K, Cs), M3AlF6 (M ¼ Li, Na, K), and Na5Al3F14 type, Solid State Sci., 8, 798–806 (2006). [56] M. Ahrens, K. Schuschke, S. Redmer and E. Kemnitz, Transparent ceramics from sol-gel derived elpasolites by cold pressing, Solid State Sci., 9, 833–837 (2007). [57] St. Wuttke, S. M. Coman, G. Scholz, H. Kirmse, A. Vimont, M. Daturi, S. L. M. Schroeder and E. Kemnitz, Novel Sol-Gel Synthesis of acidic MgF2-x(OH)x materials, Chem. Eur. J. 14, 11488–11499 (2008). [58] S. M. Coman, S. Wuttke, A. Vimont, M. Daturi and E. Kemnitz, Catalytic performance of nanoscopic - AlF3 based catalysts in the synthesis of (all-rac)--tocopherol, Adv. Synth. Catal., 350, 2517–2524 (2008).

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[59] S. M. Coman, P. Patil, St. Wuttke and E. Kemnitz, Cyclisation of citronellal over heterogeneous inorganic fluoride – highly chemo- and diasterioselective catalysts for (–)-isopulegol, Chem. Commun. 460–462 (2009). [60] K. Scheurell and E. Kemnitz, Amorphous aluminium fluoride as new matrix for vanadium containing catalysts, J. Mater. Chem., 15, 4845–4853 (2005). [61] Pratap T. Patil, A. Dimitrov, J. Radnik and E. Kemnitz, Sol-gel synthesis of metal fluorides supported Pd catalysts for Suzuki coupling, J. Mater. Chem., 18, 1632–1635 (2008). [62] G. Li, L. Wang, H. Ni and C. U. Pittmann, Jr., Polyhedral Oligomeric Silsesquioxane (POSS) Polymers and Copolymers: A Review, J. Inorg. Organomet. Polym. 11, 123–154 (2002). [63] S. Ku¨hl, Organisch modifizierte Metallfluoride – anorganisch-organische Hybridsysteme durch Sol-Gel-Synthese, diploma thesis, Humboldt University of Berlin, (2007). [64] G. Eltanany, St. Rudiger and E. Kemnitz, Supported high surface AlF3: a very strong solid Lewis acid for catalytic applications, J. Mater. Chem., 18, 2268–2275 (2008). [65] T. Murata, H. Ishizawa, I. Motoyama and A. Tanaka, Praparation of high-performance optical coatings with fluoride nanoparticle films made from autoclaved sols, Applied Optics 45, 1465–1468 (2006). [66] H. Kru¨ger, E. Kemnitz, A. Hertwig and U. Beck, Transparent MgF2-films by sol-gel coating: synthesis and optical properties, Thin Solid Films 516, 4175–4177 (2008). [67] H. Kru¨ger, E. Kemnitz, A. Hertwig and U. Beck, Moderate temperature sol-gel deposition of magnesium fluoride films for optical applications: A study on homogeneity using spectroscopic ellipsometry, Phys. Stat. Sol. 205, 821–824 (2008).

2 Microwave-Assisted Route Towards Fluorinated Nanomaterials Damien Dambournet, Alain Demourgues and Alain Tressaud Institut de Chimie de la Matie`re Condense´e de Bordeaux (ICMCB-CNRS), Universite´ Bordeaux 1. 87 Avenue du Dr Albert Schweitzer, 33608 Pessac Cedex, France

2.1

Introduction

The development of synthesis routes for the preparation of new inorganic fluorides is of continuing interest. There are numerous routes for synthesizing metal fluorides that include solid-state reactions, aqueous and sol-gel synthesis, gas-phase reactions and decomposition reactions.1,2 Since the beginning of the millennium, the synthesis of nanomaterials has gained in importance due to the intrinsic size-dependent properties of the resulting solids. In the field of synthesis techniques, the use of microwaves as a heating source is a growing area due to its peculiar heating mode.3 In the traditional heating mode, heat is transferred from the vessel toward the mixture. Microwave heating, however, takes place directly in the core of the matrix, allowing shorter reaction times and the absence of temperature gradient in the medium. It also offers the possibility of stabilizing new nanosized materials in an energy-efficient way. A large range of materials has been successfully prepared including oxides,4 phosphates,5 sulphides,6 hybrid fluorides,7 metals,8 nanoporous materials9 . . . However, the application of such a technique for preparing nanosized metal fluorides is very limited so far.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Edited by Alain Tressaud 2010 John Wiley & Sons, Ltd

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Functionalized Inorganic Fluorides

The first part of this chapter undertakes an introduction to microwave synthesis. After presenting the origin of this technique, some basic concepts of microwave heating are described and the advantages offered by this route are briefly discussed. Finally, a typical microwave oven dedicated to the synthesis is shown. Some examples of nanosized metal fluorides prepared by microwave irradiation are presented in the subsequent section. Particular attention is given to the preparation of nanosized aluminium-based fluoride materials. The impact of some synthesis parameters is discussed and a detailed characterization of the Al-based compounds is proposed to highlight the potential of this preparation route.

2.2 2.2.1

Introduction to Microwave Synthesis A Brief History

The potential use of microwave irradiation as a heating source was unexpectedly discovered by Percy L. Spencer. In 1945, while working on the development of radar through the use of a magnetron, he discovered that a chocolate bar in his pocket had melted. Thereafter, he confirmed his discovery and proposed through a patent the use of the microwave irradiation to heat some foodstuffs.10 In the 1970s, the development of this technology led to the largescale use of domestic microwave ovens. In 1986, two papers11 reported that several organic reactions could be accelerated using microwave irradiation. Since then, the use of microwave as a heating source in chemistry has continued to grow regularly. Nevertheless, at an early stage of the method, the use of the domestic microwave oven without any temperature/ pressure controls has led to unsafe and irreproducible works.12 Fortunately, the growth of this technology has enabled, more recently, the development of microwave oven dedicated to the chemical synthesis, leading to safer works and reproducible results.

2.2.2

Mechanisms to Generate Heat

Microwaves are electromagnetic radiations lying between radio wave frequencies and infrared frequencies (between 0.3 and 300 GHz). These are produced by a magnetron, which consists of a thermionic diode having an anode and a directly heated cathode.13 Microwaves contain an electric and a magnetic field component. It is the interaction between the electric field component and the matter that generates the heat through two mechanisms.13 (i) Dipolar polarization. A molecule that possesses a dipolar moment, such as water, is sensitive to an electric field. Under the influence of the latter, the dipole will attempt to align itself by rotation. The applied field provides the energy for this rotation, which depends on the frequency. The rapid direction change of the microwave components cannot be followed by the dipole. Therefore, a phase difference between the orientation of the field and that of the dipole is generated. This phase difference creates energy, which is lost from the dipole by molecular friction and collisions, giving rise to heat (dielectric heating).

Microwave-Assisted Route Towards Fluorinated Nanomaterials

41

The frequencies that allow microwave dielectric heating to take place are 918 MHz and 2.45 GHz, the latter being the most used. This mechanism depends on the ability of the dipole to reorientate under the applied electric field. The capability of a substance to convert the electromagnetic energy into heat is given by the dielectric loss: tan that is equal to the ratio of the dielectric loss "00 to the relative permittivity "0 . The relative permittivity represents the measure of the ability of a molecule to be polarized, while the dielectric loss is the ability of a medium to convert dielectric energy into heat. (ii) Ionic conduction. In a solution, the ions will move under the influence of an electric field component. As a consequence, the collision rate increases, converting the kinetic energy into heat. This phenomenon explains why the temperature of two samples containing distilled water and tap water will be higher in the case of tap water after similar microwave irradiation conditions.14

2.2.3

Advantages of Microwave Synthesis

Microwave synthesis offers a large range of advantages over conventional synthesis. The consequences of the direct heating are the following: reduction of energy consumption, better yield and gain in the reaction duration. From a practical viewpoint, the shorter reaction time enables a faster optimization of the synthesis parameters. The ‘microwave chemistry’ can be applied to practically all inorganic families that do exist (see Introduction). Based on the literature, some interesting facts have been reported and can be summarized as follows: stabilization of metastable and novel phases, increase of the phase purity and phase selectivity, short crystallization times and narrow particle-size distribution. The explanations for these advantages are still poorly understood but several reasons9 are usually pointed out: very rapid heating rate, uniform heating, unusual interactions between species in the reaction mixture, occurrence of hot spots and enhancement of the dissolution of the reacting species. The occurrence of specific microwave effects is still under debate to explain the whole advantages displayed by this technique. For instance, an hypothesis based on enhanced diffusion mechanisms has been recently proposed.15 Some breakthroughs should be expected to occur with the development of in situ techniques.9

2.2.4

Examples of Microwave Experiments

As mentioned above, the development of microwave-oven engineering has allowed a drastic enhancement of safety and reproducibility. Several ovens dedicated to chemistry are now available. For further details, the reader can referred to the paper from Kremsner et al.16 In the second part of this chapter, different compounds have been synthesized using an advanced microwave oven from the CEM Corporation (Figure 2.1). The oven is a MARS-5 Microwave Digestion System operating at a frequency of 2.45 GHz and a maximum power of 1200 W. The reactor used (XP-1500 plus model) can operate up to 220 C and 55 bars as internal pressure. To prevent explosion, any deviation in the predetermined parameters causes the system to stop. Secondly, in case of overpressure, the breaking of a disk occurs and the released gaseous species are evacuated by the

42

Functionalized Inorganic Fluorides

ventilation system. The temperature is regulated by percentage increments of the microwave power and controlled by an optical fiber. Internal pressure is measured by a pressure sensor. Several parameters can be tuned: temperature, pressure, time, heating rate and microwave power. In a typical synthesis procedure, the precursor solution is placed in a Teflon container that is fixed onto a rotating device (Figure 2.1, left). After cooling down to room temperature, the reactor is open and set in a second device that allows the solvent evaporation (right). This final step is performed under primary vacuum and argon flow.

Figure 2.1 Images of a microwave oven dedicated to chemical reactions

2.3 2.3.1

Preparation of Nanosized Metal Fluorides Aluminium-based Fluoride Materials

Numerous studies have been devoted to the investigation of aluminium fluorides. Owing to the affinity17 of fluoride ions toward Al3þ and the strong electronegativity of F ions, the Al-F system gives rise to numerous crystallographic forms.18–23 In a general way, aluminium fluorides are obtained by thermal decomposition of a precursor, which determines the final structural arrangement.24 Scheme 2.1 summarizes the synthesis of the various forms of AlF3 and their transformation toward the thermodynamically stable -AlF3 variety (ReO3-derived type).25 Aluminium fluorides were reported to be particularly suitable catalysts in the synthesis of CFCs substitutes. Recall that, owing to the strong stability of the F–C bond as compared to those of fluorine with other elements (N, Cl, Br, H . . .)24 organic fluorine chemistry has been strongly developed,26, 27 leading, for instance, to the worldwide use of chlorofluorocarbons (CFCs) in industry (refrigerants, aerosol, solvents . . .). The role of these compounds in the ozone-depleting phenomenon28 led to a drastic rise in the research into heterogeneous catalysts that would be able to convert the CFCs into more environmentally acceptable compounds – hydrogenfluorocarbons

Microwave-Assisted Route Towards Fluorinated Nanomaterials

43

(HFCs).29 In the Al-F system, the hexagonal tungsten bronze type structure, labelled HTB or -AlF3, is the most known compound because of its high catalytic properties toward halogen exchange reaction.30 Beside efforts to understand their catalytic activity, synthetic routes were developed leading to new metastable forms of AlF3. Herron et al.23, 31 developed a nonaqueous route to prepare fluoroaluminate salts with general formula MAlF4, where M is an organic cation (M ¼ pyridineHþ, N(CH3)4þ, NH4þ). The thermal decomposition of these salts leads to new crystalline AlF3 through the release of MF gaseous species (Scheme 2.1). As far as the catalytic properties are concerned, the reactivity of the aluminium fluorides is also governed by the specific surface area displayed by the solid. Because of the strong electronegativity and reactivity of fluoride ions, conventional syntheses of AlF3 result in low surface area (<50 m2.g1) materials, thus limiting their reactivity. To overcome this sintering effect, new routes have been developed using either the plasma fluorination process of a Si-riched zeolithe,32 or a nonaqueous sol-gel33 synthesis. Both routes lead to X-ray amorphous aluminium fluorides having very high surface area – i.e. 190 m2.g1 for the plasma and higher than 200 m2.g1 for the sol-gel process, and very strong Lewis acidity.34 Nevertheless, the X-ray amorphous features of these solids render more difficult the understanding of their reactivity, which has to be correlated with the local structural features. Concomitantly, the development of ab initio calculations35–37 to model the surface of crystalline AlF3 is a promising tool to understand the relationships between the reactivity and the structural modifications occurring at the surface of the crystallite. Additionally, using molecular dynamic simulation coupled with atomic pair distribution function analysis,38 it was shown that strong structural rearrangement occurs at the surface of nanoparticles with the identification of some structural motifs of -AlF3 at the surface of nano-sized -AlF3. Beside the size effect, the chemical composition of the solid is also of great importance regarding its reactivity. The competition between fluoride and hydroxyl species39 that takes place during the synthesis of these compounds offers the possibility of tuning the acidity of the solid. The incorporation of OH groups is assumed to decrease the Lewis acidity while a Brønsted one is created. By tuning the F/OH molar ratio, a large acidic scale is therefore achievable and can lead to various reactivities and catalytic activities. For instance, it has been proved that alkylation reaction was promoted by Brønsted acid sites present at the surface of a fluorinated alumina.40

α-NH4AlF4 α-AlF3.3H2O [C5H5NH] AlF4 β-NH4AlF4 [N(CH3)4]AlF4.H2O

Δ Δ Δ Δ

β-AlF3 η-AlF3 κ-AlF3 θ-AlF3

Δ

α-AlF3 (R-3c)

Δ cubic-AlF3 (Pm-3m)

Δ β-AlF3.3H2O

Scheme 2.1 Summary of the synthesis of the various AlF3 varieties prepared by thermal decomposition processes

44

Functionalized Inorganic Fluorides

In this section, a description of the use of the microwave-assisted synthesis for the preparation of nanosized and crystalline aluminium-based fluorides is proposed. A detailed characterization of the solid obtained is also presented. Using this method, three different crystalline forms can be obtained: (i) an aluminium (hydroxyl)-fluoride adopting the HTB type structure and labelled as -AlF3,41 (ii) a hydroxyfluoride adopting the pyrochlore-type structure42 and a new aluminium fluoride hydrate43 exhibiting a ReO3–derived structure. Table 2.1 summarizes the various synthesis conditions to obtain the above phases.

Table 2.1 Structural type, chemical composition, synthesis conditions of various Al-based fluorocompounds prepared by microwave irradiation Al3þ precursor

Solvents

Synthesis duration/ temp.

2

Isopropoxyde

1 h 160 C

AlF2,6(OH)0,4

3

Nitrate

Al0.82&0.18F2.46(H2O)0.54

3.6

Chloride

Water isopropanol ether Water isopropanol Water isopropanol

Structural type

Chemical composition

Pyrochlore

AlF1.7(OH)1.3

HTB ReO3

HF/Al molar ratio

2 h 160 C 1 h 140 C

Some synthesis parameters have been proved to be decisive for achieving both the chosen final form and nanosized crystallites. The following part will show the impact of the various synthesis parameters on the phase purity, crystal structure and surface areas of the obtained materials. 2.3.1.1

Synthesis Features: Impact of some Parameters

2.3.1.1.1 Effect of the [HF]/[Al] Molar Ratio The variation of the [HF]/[Al] molar ratio has been proved to be a relevant parameter with respect to the final product. Figure 2.2 shows the powder X-ray diffraction patterns obtained from different R ¼ [HF]/[Al] molar ratios. Three domains can be evidenced. For low HF contents, corresponding to R ¼ 2, the aluminium hydroxyfluoride exhibiting the pyrochlore-type20 structure is obtained. An increase of the molar ratio up to 2.8 leads to a phase mixture containing the pyrochlore and the HTB type structure. The latter is obtained as a single phase for R ¼ 3. Finally, for R > 3, besides the HTB phase, the thermodynamically stable phase19 -AlF3 appears and became predominant with increasing HF contents. Nevertheless, in these conditions the -AlF3 phase cannot be prepared as a single phase, the HTB phase being always detected even for large HF contents. 2.3.1.1.2 Effect of the Nature of the Aluminium Precursor The nature of the aluminium precursor has been proved to affect both the nature of the stabilized phase and its morphology.

Microwave-Assisted Route Towards Fluorinated Nanomaterials

45

R = 3.5 HTB + α-AlF3

R=3 HTB

R = 2.8 Pyrochlore + HTB

R=2 Pyrochlore 10

20

30

40 50 2θ (°)

60

70

80

Figure 2.2 Effect of the R ¼ [HF]/[Al] molar ratio on the X-ray diffraction powder pattern. For each X-ray diagram, the R molar ratio is noticed as well as the stabilized phase. Syntheses were conducted at T ¼ 160 C, t ¼ 2 h using Al(NO3)3.9 H2O and water/isopropanol as solvents

Influence on the Nature of the Stabilized Form For R ¼ 3, the use of an Al nitrate precursor favours the formation of the HTB phase while the use of chloride leads to a phase mixture containing -AlF3 as the major phase.41 The stabilization of the HTB phase is in fact related to a microwave-induced side reaction. There are several reports on the ability of microwave irradiation to induce side reactions such as condensation of alcohol molecules,44 decomposition of organic molecules45 and ionic liquids.46 When using nitrate as precursor, an exothermic effect is observed during the synthesis leading a drastic increase of both pressure and temperature. Such a reaction implies a redox process identified by chemical analysis, as the reduction of nitrate into ammonium ions coupled with the oxidation of isopropanol into ketones. The occurrence of NH4þ in the solid can be identified by FTIR spectroscopy. An increase in the synthesis temperature up to 170 C also confirmed the reduction of nitrate through the XRD identification of -NH4AlF4 as a side product. By analogy with the conventional route to HTB compounds, that is the decomposition of an hydrate or ammonium fluoride salts,39 it can be suggested that the ammonium ions act as a template for the stabilization of the phase. Another example of the effect of the cationic precursor on the stabilized form is displayed in Figure 2.3. For high HF contents (R ¼ 3.5), the use of nitrate as aluminium precursor leads to the stabilization of a phase mixture. In similar conditions, R ¼ 3.5, the use of aluminium chloride enables to get a single phase which XRD diagram is related to the ReO3 type structure (Figure 2.3). Influence on the Morphology The synthesis of the pyrochlore hydroxyfluoride using nitrate as precursor leads to the preparation of very well crystallized compounds with the drawback of a very low specific surface area (6 m2.g1). By replacing nitrate by

46

Functionalized Inorganic Fluorides

AlCl3.9H2O ReO3 Phase type

Al(NO3)3.9H2O HTB +α-AlF3

10

20

30

40

50

60

70

80

2θ (°)

Figure 2.3 Influence of the aluminium precursor on the X-ray diffraction powder pattern for R ¼ 3.5. For each X-ray diagram, the precursor and the final form are noted

isopropoxide, a broadening of the X-ray peaks is observed (Figure 2.4) suggesting a lowering of the crystallite size. This trend is confirmed by the increase of the specific surface area, from 6 to 77 m2.g1. The crystallization growth of pyrochlore crystallites is thus strongly affected by the nature the aluminium precursor: the weakening of the crystallization growth could be due to the steric hindrance of the isopropoxide ligand. 2.3.1.1.3 Effect of the Solvents The nature of the solvents is very important, affecting both the morphology of the solid and its chemical composition. At first, two examples highlight the impact of the solvent when

Al(OiPr)3 S = 77 m2.g−1

Al(NO3)3.9H2O S = 6 m2.g−1

10

20

30

40

50

60

70

80

2θ (°)

Figure 2.4 Dependence of the X-ray diffraction powder pattern of the pyrochlore hydroxyfluoride on the nature of the aluminium precursor. For each X-ray diagram, the precursor and the specific surface area is noted

Microwave-Assisted Route Towards Fluorinated Nanomaterials

47

targeting high surface area materials. In the case of the pyrochlore hydroxyfluoride, the addition of a small amount of ether drastically improves the surface area of the solid, from 77 to 140 m2.g1.42 Following the work of Kemnitz et al.,33 ether molecule in the presence of HF has been suggested to form an oxonium ion giving rise to ion pairs with F. Such a complex should lead to a decrease of the fluorine reactivity known to be a structuring agent and therefore enable the preparation of nanosized crystallites. During the synthesis of the HTB-type compound, the surface area could be monitored by tuning the volume ratio V ¼ Vwater/Visopropanol of the used solvents. For V ¼ 1, a surface area of 82 m2.g1 is obtained, whereas for V ¼ 6, the value drops to 3 m2.g1. Interestingly, both compounds not only differ in their morphology but also by their chemical composition. Using elemental analysis and high field 27Al NMR spectroscopy, the chemical composition of the low and high surface area compounds has been accurately determined as AlF2.4(OH)0.6 and AlF2.6(OH)0.4, respectively. Water thus favours the stabilization of OH groups inside the structure while isopropanol enables the stabilization of F ions. This trend is confirmed when using V>1, the powder X-ray diffraction pattern of the corresponding solid displaying some traces of -AlF3, that is a more fluorinated compound. 2.3.1.1.4 Coupling Sol-gel Alkoxy-fluoride Route and Microwave Irradiation A more sophisticated route has been proposed by combining nonaqueous sol-gel fluoride synthesis and microwave solvothermal process.47 The nonaqueous sol-gel process developed by Kemnitz et al.,33 enables the preparation of X-ray amorphous metal fluorides through two steps. First, an alcoholic solution of a metal alkoxide is partially fluorinated by anhydrous hydrogen fluoride previously dissolved in an organic solvent through the reaction: M(OR)3 þ xHF ! MFx(OR)3-x þ xROH.48 After drying, the resulting alkoxy-fluoride is fluorinated using gaseous CFCs or anhydrous HF, leading to X-ray amorphous metal fluorides with very high surface area. In a second step, the alkoxy-fluoride sol-gel is subjected to a microwave solvothermal process at various temperatures ranging from 90 to 200 C for 1 hour treatment. Finally, the mixture is dried under microwave irradiation but this step has no impact on the morphology of the final compound. The characteristics of the dry gel obtained by microwave synthesis are gathered in Table 2.3. While low-temperature microwave treatment (<180 C) leads to X-ray amorphous alkoxy-fluorides, a crystallization process starts at T ¼ 180 C with the detection of -AlF3 and an unidentified phase. Such an evolution is clearly related to an increase of the kinetics of fluorination. At this stage, the resulting compounds exhibit extremely high surface areas (Table 2.3) reaching 525 m2.g1 for the dry gel obtained at 180 C. At a higher temperature, 200 C, there is clear evidence for the occurrence of a complete crystallization process through: (i) the XRD identification of -AlF3 (HTB), (ii) a large increase of the fluoride ions substituting the alkoxy ligands and accounting for the very low carbon content and (iii) a decrease of the surface area dropping from 525 to 125 m2.g1.A strong increase in the internal pressure also occurs during the microwave treatment at 200 C and was suggested to arise from the breaking of the Al-OiPr bonds through the following mechanism: Al–O–CH–(CH3)2 ! Al–OH þ CH2¼CH–CH3. Such a reaction has been proposed by analogy with the behaviour of an alkoxy-fluoride upon fluorination.49 Another side reaction has been suggested to occur, based on the work from K. G.

48

Functionalized Inorganic Fluorides

Rao et al.,44 which reported the formation of ether and water molecules through alcohol condensation reaction. As reported above, the stabilization of the HTB metastable phase also takes place through the occurrence of a side reaction.

Table 2.2

Textural properties of the Al-based materials prepared by microwave irradiation

Structural type

Chemical composition

Micro-strains (Dd/d)

Pyrochlore HTB ReO3

AlF1.7(OH)1.3 AlF2,6(OH)0,4 Al0.82&0.18F2.46(H2O)0.54

2.7.103 2.2.103 2.103

Coherent domains (nm) 12 15 50

Surface area (m2.g1) 137 82 61

Table 2.3 Characteristics of the dry gel obtained by microwave treatment of an alkoxi-fluoride gel Synthesis temperature (C)

XRD

Elemental analysis Carbon content (wt %)

90 130 180 200

Amorphous Amorphous -AlF3 þunidentified phase -AlF3

Surface area (m2.g1)

F/Al

n.d n.d 12

1.5 1.7 2

425 470 525

0.1

2.7

125

The multi-component compound prepared at 180 C is of interest because of its very high surface area (525 m2.g1). This property is in contrast with the detection of highly crystallized X-ray peaks identified as -AlF3, thus suggesting that the solid is mainly built from small particles, which should contain some organics moieties. These species can be successfully removed by low temperature fluorination using pure F2 gas at 225 C preserving a high surface area of 330 m2.g1. XRD analysis confirmed the thermal stability of both -AlF3 and the unidentified phase. HRTEM characterization of the multi-component compound revealed that the solid is build of large particles (50 nm) identified as -AlF3 and small particles (10 nm and less) related to the second component. An example of such small crystallized domains is displayed in Figure 2.5. A zoom on one particle (inset Figure 2.5) shows interplanar distances of about 0.30 nm, which is very close to the one detected by XRD (0.31 nm), thus confirming that the unidentified phase consists of very small particles and is the main constituent of this compound. Finally, this multi-component material displays high catalytic properties toward halogen exchange reactions that can be correlated to the large number of strong Lewis acid sites detected on the surface of the material.47

Microwave-Assisted Route Towards Fluorinated Nanomaterials

49

Figure 2.5 HRTEM image of HSA AlF3 (330 m2.g1) showing the main (unidentified) constituent of the solid (zones 2). Inset: zoom on a small crystalline domain exhibiting interplanar distances at 0.30 nm. (Reprinted with permission from [47] Copyright (2008) Wiley-VCH.)

2.3.1.2

Characterization of the Crystalline Al-based Materials Obtained by Microwave Irradiation

2.3.1.2.1 Textural Properties Analysis of the X-ray Line Broadening It has been shown that microwave synthesis enables the stabilization of high surface materials. The knowledge of the textural properties of the prepared compounds is a relevant point because it strongly affects their physical-chemical properties. Beside the measurement of the surface area determined by the BET method, it is important to gain some information on the microstructure of the solid – that is the coherence domains and the occurrence of micro-strains (Dd/d) inside the solid. A relevant method to determine the microstructure of a solid is the global profile refinement. A solid made of nanosized particles is always characterized by a broadening of the X-ray peaks. Such a phenomenon can be used to accurately calculate the values of the micro-strains and the crystallite size.50 The X-ray line broadening analysis is performed by fitting the X-ray peak profile by using the pseudo-Voigt function of Thompson- Cox-Hastings.51 Such a function allows the angular dependence of both Lorentzian (HL) and Gaussian (HG) components of the Full Width at Half Maximum (FWHM) to be refined separately. The instrumental contribution of the X-ray line broadening was previously determined using the XRD pattern of a reference sample (e.g. LaB6), which does not contribute to the enlargement of the X-ray peaks. The sample contribution is therefore calculated as follows: Gsize iso HG2 ¼ Ustrain iso þ ð1 Þ 2 D2strain aniso tg2 þ cos 2 HL ¼ ðXstrain

iso

þ Dstrain

aniso Þtg

þ

Ysize

iso

þ Fsize cos

aniso

where Ustrain_iso, Xstrain_iso, Gsize_iso, Ysize_iso are refinable parameters relative to isotropic strain and size broadening effects, whereas Dstrain_aniso and Fsize_anizo are analytical

50

Functionalized Inorganic Fluorides

functions for hkl-dependent model. is a mixing coefficient for the Lorentzian part of the strain. The integral method is finally applied to calculate the apparent crystallite size and strains of the sample. The corresponding values are given in Table 2.2. The specific surface area is also noted. The determination of the coherence domains confirms that the prepared solids are built of nanosized particles in agreement with large surface area. It is interesting to note that the higher the fluorine content, the larger the particle size, highlighting the structural properties of the fluoride species. Moreover, microstrains can be also induced by a nanosized effect or by the occurrence of mixed anions inside the structure. The particle size of the materials can be also studied through high-resolution transmission electron microscopy (HRTEM). Some micrographs are displayed in Figure 2.6,

A

B

C

Figure 2.6 HRTEM micrographs of isolated particles of the Al-based fluorides materials: A. Pyrochlore, B. HTB and C. ReO3 derived form. (Reprinted with permission from [41–43] A. Copyright (2008) Royal Chemical Society, and B and C, Copyright (2008) American Chemical Society.)

Microwave-Assisted Route Towards Fluorinated Nanomaterials

51

illustrating that the observed dimensions of the crystallites are in good agreement with the coherence domains calculated by X-ray line broadening. Nanostructured HTB Phase As mentioned above, the surface area of the solid can be monitored by tuning the volume ratio V ¼ Vwater/Visopropanol. Three different types of samples can be obtained with various surface areas, that is, 3, 49 and 82 m2.g1 and labelled as LSA (low surface area), MSA (medium) and HSA (high).41 When getting towards high surface area, some X-ray peaks display some anisotropy with regards to their FWHM. The profile matching of the HSA compound evidences that the FWHM of some experimental peaks was smaller than the calculated one (Figure 2.7, left). This hkl-dependent line broadening has been proved to arise from size effect and has been solved by the use of anisotropic size broadening model, that is, a platelet-shape along the a axis (Figure 2.7, right). The calculated dimensions of the platelet are 8 nm along the a-axis and 34 nm in the b, c plane. Such a model has been confirmed by HRTEM observations, with an average particle size of 15 nm (cf. Figure 2.6). The knowledge of the morphology of a crystallite is of great importance because it enables a map of the surface to be drawn up. The above result suggests that the {100} surface is weakly exposed, as compared to the {001}, {011} and {010} surfaces. Wander et al.52 have studied the equilibrium crystal morphology of AlF3 and their results predicted that the exposed crystal surfaces are: 59 % for {010}, 38 % for {001} and 3 % for {100}, proving that the last one is not a thermodynamically stable surface. Although the metastability of the {100} surface is confirmed by the above microscopy observations, the relative contribution of the {100} surface is higher than those predicted by computational data. Since the anisotropic size effect appears to be related to the nano-size, it can be suggested that the relative contribution of the {100} surface type increases with decreasing the crystallite size. Such a trend should impact on the acidity of the solid because ab initio calculations of NH3 adsorption on the (100) surface of -AlF3 reveal that this surface displays strong Lewis acidity.53 This surface also displays some high catalytic properties toward chlorine–fluorine exchange reactions.54 2.3.1.2.2 Structural Features The above sections have presented the synthesis and the textural properties of Al-based materials prepared by microwave treatment. In the following, the structural features of the three compounds will be presented in more details. Comparison between the HTB and the Pyrochlore Hydroxyfluorides Description of the networks. The pyrochlore and the hexagonal tungsten bronze frameworks are built of corner-sharing AlF6x(OH)x octahedra exhibiting rather similar arrangements. While the pyrochlore-type structure displays a three dimensional hexagonal tunnel system, the HTB form is characterized by a one-dimensional tunnel along the c-axis (Figure 2.8). The pyrochlore structure20 crystallizes in the cubic symmetry (Fd-3 m) leading to a random distribution of OH/F anions in the 48 f sites. The HTB form21 crystallizes in an orthorhombic system (Cmcm) with two unequivalent Al atoms and four unequivalent F atoms (Figure 2.8A). Aluminium-based fluoride structures can be described by the connection of building units5 (Table 2.4). The pyrochlore network can be drawn through the linkage of three- and six-membered rings, the latter forming the 3 D interconnected channel system. In the HTB form, anionic crystallographic sites, labelled F3 and F4, allow connection of the planes containing the three- and six-membered rings resulting in a four-membered ring generating the one-dimensional hexagonal channel system along the c axis (Figure 2.8B).

52

Functionalized Inorganic Fluorides 11 000

7700

002

6700

9000

4700

Intensity (arb. units)

Intensity (arb. units)

5700

3700 2700 1700 700

7000 5000 3000 1000

–300 –1000

–1300 –2300 22.6 23.1 23.6 24.1 24.6 25.1 25.6 26.1 26.6 27.1 27.6 2θ (°)

–3000 10

20

30

40

50

60 70 2θ (°)

80

90

100 110

Figure 2.7 Zoom on the 002 line of the powder X-ray diffraction pattern of the HTB phase, illustrating the anisotropy of the X-ray peaks (left). Final profile matching obtained with the use of a platelet shape model. Dotted line: experimental, full line: calculated (right)

b A F3

a

F4 F2

F1

B c b

C b

a c

Figure 2.8 b-AlF3 structure (HTB-type) along (A) the c-axis, (B) the a-axis. The various anionic crystallographic sites have been labelled. The Al(1) and Al(2) atom type are represented as light spheres. (C) Pyrochlore network

Microwave-Assisted Route Towards Fluorinated Nanomaterials Table 2.4

Structural characteristics of a-, HTB- and pyrochlore forms of AlF3

Phase

Building units

Al-X-Al (X¼ F or OH) angles ()

-AlF3 -AlF3 Pyrochlore

4 3–6–4 3–6

158 148– 148 – 166 141

Table 2.5

Average interatomic ˚) distances (A 1.797 1.800 1.842

Crystallographic data obtained by Rietveld refinement

Crystal symmetry Space Group Z Unit cell ˚) parameters (A ˚ 3) Volume (A ˚) Interatomic distances(A

Angles ()

53

-AlF2.6(OH)0.4 (HTB)

Pyrochlore AlF1.7(OH)1.3

Orthorhombic Cmcm 12 6.9681(2) 12.0360(3) 7.1434(1) 599 Al2-F1 2 1.857(12) Al1/Al2 - F2 4 1.77(3)/2 1.82(2) Al1-F3 2 1.789(5) Al2-F4 2 1.804(8) Al2-F1-Al2 139.4(5) Al1-F2-Al1 151.9(11) Al1-F3-Al1 173.1(2) Al2-F4-Al2 163.8(4)

Cubic Fd-3 m 16 9.7309(1) 921 1.824(1)

141.12(6)

Through the effect of the [HF]/[Al] molar ratio, it has been shown that the competition between hydroxyl and fluoride groups has a strong influence in the stabilization of the final compound. To understand the impact of the nature of the Al-X bond (X¼OH, F) on the structure, conventional literature data (Table 2.4) are compared to those obtained by Rietveld refinement for the compounds prepared by microwave irradiation (Table 2.5). The impact of OH groups. Due to the strong electronegativity of fluoride ions, Al-F bond displays shorter interatomic distance than Al-O(-OH) ones. This is well evidenced by the bond distances of the different AlF3 forms: 1.797 for -AlF3, 1.80 for -AlF3 whereas the hydroxyfluoride pyrochlore type structure, AlF1.5(OH)1.5, exhibits Al-X distances of ˚ . A slight variation of the chemical composition is sufficient to decrease the 1.842 A interatomic distances because the pyrochlore with a chemical composition AlF1.7(OH)1.3 ˚ . The occurrence of OH groups has been also suggest to exhibits distances of 1.824 A favour small Al-X-Al angles as typically found in the three and six- membered rings.55 The impact of the OH groups inside the framework is relevant in the case of the HTB form. Using conventional route, the HTB form contains some traces of OH groups while the microwave synthesis leads to the preparation of an hydroxyfluoride phase. The comparison between both compounds clearly highlights the impact of the OH groups upon the structure. By comparison with the hydroxyl-free compound -AlF3, it has been shown that

54

Functionalized Inorganic Fluorides

F1 and F2 sites are partially hydroxylated. At first, the Al-F1 bond is characterized by an ˚ , significantly larger than those found in the hydroxyl-free interatomic distance of 1.857 A 41 ˚ compound that is 1.800 A. The value of the Al-F1-Al angle (141) is also close to those found in the pyrochlore type structure, confirming the partial hydroxylation that occurs in this site and the trend of OH groups to form smaller Al-X-Al angle. Concerning the F2 site, the Al-F2 interatomic distances display a large estimated standard deviation, which also suggests the occurrence of OH groups in this site. On the other hand, the F3 and F4 crystallographic sites exhibit typical Al-F interatomic distances revealing the absence of hydroxyl groups in these sites. From a structural point of view, the hydroxylation of the HTB-type structure occurs in the anionic sites that form the structural pattern that is threeand six-membered rings of the pyrochlore. This is in agreement with the observation that the increase of the [HF] content leads to the following sentence: pyrochlore (3 and 6 membered rings) ! HTB (three-, six- and four-membered rings) ! -AlF3 (fourmembered rings) (see Table 2.4) and finally clearly highlight the impact of the OH groups upon the structure.41 While the structural relationship between the HTB and the pyrochlore phases has been understood, it appears interesting to move from long to short range order with the use of more local probe such as the NMR and XPS spectroscopy. Local environments. 19F and 27Al NMR spectroscopy is a powerful tool to probe the anionic repartition as well as the cationic disorder. 19F NMR gives some precise information about the OH/F repartition because it is well known that the 19F chemical shifts of octahedral AlF6x(OH)x environments increase with the oxygen content. The 19F chemical shifts for the pyrochlore and the HTB type structure obtained from spectrum reconstruction are gathered in Table 2.6. The 19F MAS NMR spectrum of the pyrochlore presents only one contribution, whereas the 19F MAS NMR spectrum of the HTB has been reconstructed using three contributions. Note that the iso of a fluorine atom in an AlF63 octahedron is located at 172 ppm, as found for - and -AlF3.56 The OH/F content in the two compounds is clearly evidenced through the position of the iso. Another remarkable feature in the 19F MAS NMR spectra of both compounds is the line width (Figure 2.9). In the pyrochlore, the peaks are significantly broader than those of the HTB one’s indicating a disorder around fluorine atoms in this compound. This is due to the random distribution that occurs for the OH/F atoms in the 48 f site. Concerning the aluminium environments, the use of high magnetic field (17.6 T) 27Al NMR spectroscopy allows a better resolution of the iso value by the reduction of the quadrupolar interaction contribution of the spectrum. Figure 2.10 presents the 27Al MAS

Table 2.6 Isotropic chemical shift diso and relative intensity deduced from the reconstruction of the 19F MAS NMR spectra Phase

iso (ppm)

Pyrochlore AlF1.7(OH)1..3 HTB AlF2.6(OH)0.4

161.2 173.5 167.6 166.7

Intensity (%) 100 28 59 13

Microwave-Assisted Route Towards Fluorinated Nanomaterials

0

–50

–100

–150 –200 δiso (ppm)

–250

55

–300

Figure 2.9 19F MAS NMR spectra of the pyrochlore AlF1.7(OH)1.3 (dotted line) and the b-AlF2.6(OH)0.4 HTB-type structure. (Asterisks: spinning side bands.) (Reprinted with permission from [42] Copyright (2008) Royal Society of Chemistry.)

NMR spectra of the two hydroxyfluorides and Table 2.7 gives the characteristics of the 27 Al MAS NMR spectra obtained after reconstruction. Both spectra display several contributions. As a reference, - and -AlF3 exhibit iso at –16 and –15 ppm, respectively. As for 19F MAS NMR spectroscopy, the occurrence of hydroxyls in the vicinity of Al3þ ions leads to an increase of the iso value. The difference between both spectra also reveals the higher OH/F ratio displayed by the pyrochlore. As both structures consist of corner-sharing AlF6x(OH)x octahedra, the assignments of the 27Al NMR signals can be performed by considering a distribution of OH groups in AlF6 octahedra. While the pyrochlore phase displays AlF6x(OH)x species with x ranging from 0 to 6, in HTB phase x values range from 2 to 6 because of the occurrence of pure fluorinated F3 and F4 sites. Based on this AlF6x(OH)x distribution, the probability of having various AlF6x(OH)x species can be calculated using a binomial law, which matches well with the experimental data (Table 2.7). Such results enable the identification and quantification of the various AlF6x(OH)x species forming the framework. Additionally, the impact of the OH groups on the distortion of the octahedra is clearly reflected by the Q value, which increases with increasing the number of OH groups. The electronic structure of the Al3þ ions has been also probed using X-ray photoelectron spectroscopy. The presence of a strong electroattractive anion such as fluoride ions leads to an increase of the Al 2p binding energies.57,58 The Al 2p XPS spectra are displayed in Figure 2.11 and the XPS data obtained by spectrum reconstruction are gathered in Table 2.8. The Al 2p XPS spectra have been fitted using three components. The spectrum of the pyrochlore is dominated by the component located at 76 eV and ascribed to a mixed F/OH anion environment around the Al3þ ion. Some weak contributions are also detected at low and high binding energies, which account for hydroxylated and fluorinated environments in agreement with the random distribution of OH and F anions that occurs in the 48f site. The Al 2p XPS spectrum of the HTB phase reveals a large heterogeneity in the electronic density of Al3þ ions. Contrarily to the averaged OH/F

56

Functionalized Inorganic Fluorides

environment found in the pyrochlore, the HTB phase displays various environments: hydroxylated, fluorinated and mixed OH/F. This can be rationalized by the large range of ˚ . The occurrence of distinct hydroxylated interatomic distances lying from 1.78 to 1.85 A and fluorinated sites is assumed to induce an electronic heterogeneity around the Al3þ ions. It should be noted that for both phases three components have been also detected for F1s BEs i.e. 686.8, 685.7 and 684.7 eV. Interestingly, the XPS data reflects the acidity displayed by both solids. Adsorption of probe molecules has indeed confirmed the homogeneous and heterogeneous acidity for the pyrochlore42 and HTB59 phases, respectively.

2

1

3 Pyrochlore

4 1

AlF1.7(OH)1.3

2 3 10

0

−20 −10 δiso (ppm)

β-AlF2.6(OH)0.4 −30

−40

Figure 2.10 27Al MAS NMR spectra of the pyrochlore AlF1.7(OH)1.3 and the HTB-type b-AlF2.6(OH)0.4 compounds Table 2.7 Line label, isotropic chemical shift diso (ppm), quadrupolar product Q (kHz), relative line intensity (%) deduced from the reconstruction of the 27Al MAS NMR spectrum, with corresponding assignments. For a sake of clarity, the very weak 27Al NMR signals located at diso >10 ppm ascribed to the partial decomposition of the pyrochlore has not been mentioned Line

iso (–0.5)

Q (–5)

1 2 3 4

–12.5 –8.7 –4.2 1.8

345 540 625 505

1 2 3

–15.5 –11.7 –9.5

280 610 990

Relative Intensity (–0.5)

Assignment of the 27Al NMR lines

Pyrochlore AlF1.7(OH)1.3 47.3 AlF6, AlF5(OH) and AlF4(OH)2 33.4 AlF3(OH)3 15.5 AlF2(OH)4 3.9 AlF(OH)5 -AlF2.6(OH)0.4 82 16 2

AlF6 and AlF5(OH) AlF4(OH)2 AlF3(OH)3

Microwave-Assisted Route Towards Fluorinated Nanomaterials

57

Pyrochlore AlF1.7(OH)1.3 2

3

1

β-AlF2.6(OH)0.4

84

82

80

78

76

74

72

70

68

BE (eV)

Figure 2.11 Al 2 p XPS spectra of the pyrochlore and HTB-type compounds Table 2.8 Al 2 p XPS data of the HTB and the pyrochlore phases obtained by spectrum reconstruction using FWHM of 2.1 eV. An alignment between the components of both compounds has been undertaken with 0.1 eV as a standard deviation to overcome the reference issue Line Label

1 2 3

Al 2 p position eV (–0.1)

74.4 76 77.4

Relative intensity (%) Pyrochlore

HTB

16 73.5 10.5

22 54 24

A New Form of Aluminium Fluoride Hydrate It has been noted (cf. 2.3.1.1.2) that using chloride as aluminium precursor, a single phase whose XRD is closely related to the hightemperature phase of -AlF3 could be obtained (Figure 2.12). It can be added that using a higher synthesis temperature, that is 180 C instead of 140 C, an additional X-ray peak is detected, suggesting the presence of -AlF3 (113) line (Inset Figure 2.12). This point highlights the need for a careful preparation to achieve phase purity. Using the International Centre for Diffraction Data (ICDD), the powder pattern could be indexed in a first step as a cubic phase ˚ ) with the chemical formula AlF3.H2O.60 Nevertheless, the structural (Pm-3 m, a ¼ 3.600 A description appeared unrealistic because of a poor reliability factor (RBragg ¼ 16.4 %). The structure of this new compound has recently been investigated.43 At first, FTIR

58

Functionalized Inorganic Fluorides

spectroscopy and thermogravimetric analysis confirmed that the prepared compound consists indeed of a hydrated aluminium fluoride, but the attempts to place water molecule in the conventional cubic cell led to a poor agreement factor (RBragg 18 %). Water molecules were finally considered as part of the first coordination sphere of Al3þ ions. Nevertheless, because water molecule is neutral, [AlF6x(H2O)x]x3 octahedra display a lack of electron density, which leads to consider the occurrence of cationic vacancies. The latter has been clearly demonstrated by the convergence of the reliability factor down to RBragg ¼ 4.3 % when refining the occupancy rate of the Al3þ atoms and introducing water in the anionic fluorinated site (3a). The chemical formula Al0.82&0.18F2.46(H2O)0.54 could be proposed based on elemental analysis, TGA and Rietveld refinement. The crystallographic data of Al0.82&0.18F2.46(H2O)0.54 are gathered in Table 2.9 and compared to those of -AlF3. Interestingly, the interatomic distances for both hydrate and -AlF3 are very close. While the occurrence of the water molecule should lead to an increase of the interatomic distances as found for other aluminium fluoride hydrate,61 cationic vacancies may act as a repulsive entity decreasing the Al-X distances and maintaining the cubic symmetry. Table 2.9 Crystallographic data of Al0.82&0.18F2.46(H2O)0.54, a-AlF3 and the high temperature cubic form Phase

Crystal symmetry Space Group Z

Unit cell para- Interatomic ˚) ˚ )Al-X meters (A distances(A

Al0.82&0.18F2.46(H2O)0.54 Cubic Pm-3 m 1 Rhombohedral R-3 c -AlF3 12 High temperature cubic Cubic Pm-3 m 1 form

3.6067(1) 4.9305(6) 12.4462(7) 3.58

1.8034(1) 1.797(3) 1.791

Angles () 180 157.07(7) 180

T=180°C

41

20

30

40

50

43

60

45

70

80

Figure 2.12 X-ray diffraction pattern of the aluminium fluoride hydrate obtained after microwave irradiation at T ¼ 140 C. Inset: effect of the temperature on the XRD

Microwave-Assisted Route Towards Fluorinated Nanomaterials

59

In both hydroxyfluoride and hydrate compounds, the anionic sites are statically occupied either by F, OH ions or H2O molecules, thus generating disorder. Refinement of the Debye-Weller factors enables to get information about the disordering state in the anionic sites. A comparison between the three phases is displayed in Table 2.10. The hydroxyfluoride phases clearly show the impact of the OH/F substitution on the ordering state of the anionic sites, an increase of the substitution leading to higher Debye-Weller factors. While the substitution rate (F/H2O) in the anionic site is rather low, the hydrate phase Al0.82&0.18F2.46(H2O)0.54 exhibits the highest disorder. FTIR and 1H MAS NMR spectroscopy indicate that water molecule is in average located on the 3 d Wyckoff position inducing the large thermal displacement observed. The local structure has been characterized using 19F and 27Al MAS NMR (Figure 2.13). By analogy with the isotropic chemical shift of the hydroxyl-fluorinated species found in the HTB and pyrochlore phases, the AlF6 and AlF6x(H2O)x species (x ¼ 1, 2, 3) have been identified and quantified (Table 2.11). Interestingly, two contributions, reconstructed using four signals (Table 2.11), are displayed by the 19F MAS NMR spectrum (Figure 2.13). The major one ( 88 %) lies in the range of fluoride ions as found in -AlF3 and in HTB and pyrochlore hydroxyfluorides and is therefore ascribed to bridging fluorine atoms in AlF6x(H2O)x octahedra. The minor contribution (12 %) exhibits chemical shift values as observed in AlF3.9 H2O (–149.5 ppm) and -AlF3.3 H2O (–147.9 ppm), which are both built from isolated octahedra.61 The resulting fluoride atoms are thus assigned to nonbridging fluorine atoms in AlF6x(H2O)x octahedra and therefore localized next to an aluminium vacancy. The low contribution of nonbridging fluorine atoms gives the evidence that the aluminium vacancy is mostly surrounded by water molecules. FTIR spectroscopy gives information about the conformation of water molecules inside the structure. A 2:1 complex has been detected by FTIR, which can be noted as XH-O-HX where X is a proton acceptor (X ¼ F–, H2O). By comparison with IR bands of water in an inert solvent,62 the vibration bands of OH groups of water molecules are shifted due to hydrogen bonding towards lower wavenumber that is around 3185 and 3335 cm1 for sym(OH) and asym(OH), respectively. A representation of the structure is presented in Figure 2.14. 2.3.1.3

Key Points and Concluding Remarks

The above mentioned studies on aluminium-based fluorides compounds prepared by microwave-assisted synthesis have enabled the improvement of the knowledge on these materials as well as the development of different compounds having unusual properties in

Table 2.10 Comparison of the Debye-Weller factors obtained from Rietveld refinement for the Al-based fluorides prepared by microwave irradiation ˚ 2) Biso (A

Phase

Pyrochlore AlF1.7(OH)1.3 HTB AlF2.4(OH)0.6 HTB AlF2.6(OH)0.4 Al0.82&0.18F2.46(H2O)0.54

Al3þ

X (X ¼ F, OH, H2O)

0.78(2) 0.62(3) 0.79(3) 0.6(3)

1.12(2) 1.25(4) 0.83(3) 2.1(4)

Functionalized Inorganic Fluorides

60

Table 2.11 Al0.82&0.18F2.46H2O0.54 line label, isotropic chemical shift diso (ppm), quadrupolar product Q (kHz), relative line intensity (%) deduced from the reconstruction of the 27Al MAS NMR spectrum and the corresponding assignments. Isotropic chemical shift diso and relative intensity deduced from the reconstruction of the 19F MAS NMR spectrum 27

Line 1 2 3 4

Al NMR

iso (–0.5)

Q (–5)

Relative Intensity (–0.5)

Assignment

17.9 14.4 11.7 –9.4

118 252 360 559

59.1 28.7 7.8 4.4

AlF6 AlF5(H2O) AlF4(H2O)2 AlF3(H2O)3

Relative Intensity (–0.5)

Assignment

13.4 74.3 4.4 7.9

Bridging F

19

F NMR

iso(–0.5) 173.0 169.4 155.2 152.5

Non-bridging F

1 2

3 4

−140

−160 19F δ iso

−180

−200

10

0

−10 27 Al δ iso

(ppm)

−20

−30

−40

(ppm)

2θ (°)

Figure 2.13 19F and 27Al MAS NMR spectra of Al0.82&0.18F2.46(H2O)0.54. (Reprinted with permission from [43] Copyright (2008) American Chemical Society.)

terms of morphology, composition and structure. The main highlights brought by the use of the microwave synthesis are summarized below: • A strong dependence of the [HF] content on the structural arrangement has been found to occur. Such a phenomenon has been rationalized by the structural impact of the hydroxyl groups. The coupling of sol-gel and microwave synthesis has also shown the effect of the temperature on the kinetics of fluorination. • Microwave-assisted synthesis has enabled the direct stabilization of metastable phases such as -AlF3. Additionally, it has enabled the preparation of a new phase with unusual structural features that is a derived form of -AlF3 containing cationic vacancies owing to the occurrence of water molecules in the vicinity of the cation.

Microwave-Assisted Route Towards Fluorinated Nanomaterials

61

• All the prepared materials exhibit very high surface area with nanosized crystallites as opposed to compounds prepared by conventional methods. • All the syntheses presented in this part have been obtained in one or two hours, which is unconventionally faster than comparable methods such as solvo-hydrothermal treatment. This fast and energy-efficient synthesis enabled a quicker investigation of the relevant parameters. Al B F F/OH2

A

2:1 complex

Al3+ vac

Isolated H 2O

2.55 Å

Figure 2.14 Representation of Al0.82&0.18F2.46(H2O)0.54: (A) view of the cubic cell. (B) Representation of the structure with an isolated water molecule. (Reprinted with permission from [43] Copyright (2008) American Chemical Society.)

2.3.2 2.3.2.1

Microwave-assisted Synthesis of Transition Metal Oxy-Hydroxy-Fluorides Titanium Oxy-hydroxy-fluorides

2.3.2.1.1 Influence of the [HF]/[Ti] Molar Ratio Like Al-based fluorides, the [HF] content expressed as the R ¼ [HF]/[Ti] molar ratio, is a decisive parameter for the final crystallized forms. Using oxychloride (TiOCl2.xHCl) as titanium precursor, the variation of the [HF] content leads to the stabilization of three phases (Figure 2.15): (i) for R ¼ 0.2, the anatase-type structure, (ii) for R ¼ 1.7, a new phase adopting a derived-HTB type structure and finally, (iii) for R ¼ 3 a superstructure of the ReO3 form. The two latter phases are obtained at T ¼ 90 C. On the other hand, the fluorine-doped anatase is obtained at higher temperature, that is T ¼ 150 C. The composition, structural features and UV-absorption properties of Ti oxy-hydroxy fluorides are detailed in chapter 8 of this book by A. Demourgues, L. Sronek and N. Penin.69–71 2.3.2.1.2 Impact on the Chemical Composition Here again, the use of microwave-assisted route enables the synthesis of a new series of Tibased fluoride nanomaterials at low duration time and low temperature. Beside, the report of a phase derived from the HTB-type structure, the synthesis conditions favour the competition between OH and F ions, leading to new compositions besides TiOF2, the only titanium oxyfluoride reported so far by various authors through various synthesis

62

Functionalized Inorganic Fluorides

R = 3/ReO3 form

R = 1.7/HTB form-derived

R = 0.2/F-anatase

10

20

30

40

50

60

70

80

2θ (°)

Figure 2.15 Dependence of the X-ray diffraction powder pattern on the R ¼ [HF]/[Ti] molar ratio. For each X-ray diagram, the R molar ratio is noticed as well as the nature of the stabilized phase. Syntheses were conducted at T ¼ 90 C, t ¼ 30 min (TiOCl2.xHCl in water), excepted for R ¼ 0.2: T ¼ 150 C t ¼ 30 min

approaches.63,64 On the contrary, the microwave process leads, for instance, to the stabilization of an hydroxyfluoride containing some cationic vacancies: Ti0.75F1.5(OH)1.5,65 which adopts a superstructure of the ReO3 form. Moreover, the HTB-derived form corresponds to the following formula Ti0.75O0.25(OH)1.3F1.2 with cationic vacancies and anionic nonstoichiometry, whereas the anatase form Ti0.9F0.2O1.6(OH)0.2 contains also three types of anions with cationic vacancies. 2.3.2.1.3 From Aqueous to Organic Medium Effect on the stabilized phase. The impact of the medium on the stabilization of Ti-based compounds has been investigated by tuning the solvent and cationic precursor while keeping the molar ratio R equal to 1.7. Figure 2.16 shows the X-ray diffraction powder patterns obtained using either aqueous (TiOCl2.xHCl in water) or organic media (Ti(OCH(CH3)2)4 in isopropanol). While the aqueous medium leads to a well-crystallized phase identified as a new oxy-hydroxy-fluoride, the use of isopropoxyde in isopropanol enables the preparation of fluorinated anatase. According to the width of the X-ray peaks, the solid as prepared is built of nanoparticles, which is confirmed by the HRTEM images showing crystallite dimensions of about 5 nm (Figure 2.17). Fast Fourier Transforms (FFT) of HRTEM image matched with a [111] zone axis of the I41/amd tetragonal anatase structure. Such a result is rather surprising keeping in mind the strong influence of the [HF] content on the stabilized form. The kinetics of fluorination is here clearly dependent of the nature of the ligand. 2.3.2.1.4 Tin Titanium Oxy-hydroxyfluoride Xie et al.,66 have synthesized a new tin-titanium oxy-hydroxy-fluoride Sn1.24Ti1.94O3.66(OH)1.50F1.42 exhibiting a pyrochlore-type structure. The synthesis

Microwave-Assisted Route Towards Fluorinated Nanomaterials

TiOCl2.xHCl H2O

10

20

30

50

40 2θ

60

63

Ti(OCH(CH3)2)4 Isopropanol

70

80

10

20

30

40

50

60

70

80

2θ

Figure 2.16 Dependence of the X-ray diffraction powder pattern oxy-hydroxy-fluoride on the medium used: aqueous (left), organic (right)

of

Ti-based

Figure 2.17 HRTEM image of F-anatase nanoparticules prepared using organic medium. Inset is the FFT of particle ˚ showing a [111] zone axis of the I41/amd tetragonal structure

conditions were optimized by tuning several parameters, i.e. the pH of the solution. The final solid is build of nanosized crystallites ranging from 20 to 30 nm. Interestingly, a rutile phase Sn0.39Ti0.61O2 with particle size around 100 nm is obtained by the decomposition of this oxy-hydroxy-fluoride at 800 C, which is far lower than the conventional route used generally for preparing this compound, i.e. 1450 C.

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Functionalized Inorganic Fluorides

2.3.2.2

The Use of Ionic Liquids for the Synthesis of Metal Fluorides

In this case, metal fluorides are obtained by decomposition of ionic liquid. The latter is the source of F ions. The experiments have been performed in a domestic oven (the one that you are using to cook food).46 As a general method, a mixture containing the ionic liquid BMIBF4 (1-Butyl-3-methylimidazolium tetrafluoroborate) and a nitrate metal salt is subject to microwave irradiation in a domestic oven for a very short time, ranging from 5 to 10 min. In this way, the synthesis of various metal fluorides presenting nanosized and particular morphology were reported: FeF2 (nanorods), CoF2 (aggregated needles), ZnF2 (anisotropic morphology), LaF3 (oval) and YF3 (needle-shaped). The release of fluorine ions from the solvent was claimed to arise from an hydrolysis reaction such as BF4 þ H2 O ! BF3 :H2 O þ F . 2.3.2.3

Metal Organic Frameworks

In the search of porous inorganic-organics hydrides solids (see other contributions in this volume) which display several potential applications,67 the development of new synthesis routes is of crucial importance because the conventional methods generally consist of hydro- or solvo-thermal processes that require up to several days. The cubic chromium terephthalate Cr3F(H2O)2O[C6H4(CO2)2]3 (MIL-101), for instance, has been successfully synthesized using a microwave solvothermal process.68 Owing to the faster dissolution of the precursors and/or the condensation process, the MIL-101 can be obtained even after 1 min of irradiation at 210 C. Finally, a well crystallized material is obtained for 40 min of irradiation at 210 C while the conventional method is performed at 220 C for 10 h. The very rapid synthesis of MIL-101 leads to nanosized crystals that result in an improvement of the sorption properties of benzene.

2.4

Concluding Remarks

Although microwave ‘inorganic’ chemistry has been developed only recently, such a technique has already showed extraordinary potential. Its application to the preparation of metal fluorides clearly offers numerous advantages: the direct stabilization of metastable phase, the isolation of new phases, the preparation of homogenous and nanosized particles leading to high surface area materials, the easy control of the kinetic of fluorination. There are many applications of the prepared materials, including catalysis, optics and energy storage. Moreover, the unconventional heating mode enables fast reaction to occur, which renders this route one of the most energy efficient currently available. It is clear that interest in such a technique will continue to growth in the future with a need for a deeper understanding of the microwave chemistry.

Acknowledgements The EU is gratefully acknowledged for financial support through the Sixth Framework Programme (FUNFLUOS, Contract No. NMP3-CT-2004–5005575). The

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65

fluorination experiments have been carried out with the help of Etienne Durand. Most NMR results have been obtained at Universite´ du Maine by Christophe Legein and Jean-Yves Buzare´.

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[20] J. L. Fourquet, M. Rivie`re and A. Le Bail, Crystal structure and protonic conductivity of pyrochlore phase Al2[(OH)1xFx]6.H2O and Al2[(OH)1xFx]6, Eur. J. Solid State Inorg. Chem., 25, 535 (1988). [21] A. Le Bail, C. Jacoboni, M. Leblanc, R. De Pape, H. Duroy and J. L. Fourquet, Crystal structure of the metastable form of aluminum trifluoride -AlF3 and the gallium and indium homologues, J. Solid State Chem., 77, 96 (1988). [22] A. Le Bail, J. L. Fourquet, U. Bentrup t-AlF3: Crystal structure determination from X–ray powder diffraction data. A new MX3 corner-sharing octahedra 3D network, J. Solid State Chem. 199, 151 (1992). [23] N. Herron, D. L. Thorn, R. L. Harlow, G. A. Jones, J. B. Parise, J. A. Fernandez-Baca, T. Vogt, Preparation and structural characterization of two new phases of aluminum trifluoride, Chem. Mater., 7, 75–83 (1995). [24] E. Kemnitz and D. H. Menz, Fluorinated metal fluorides oxides and metal fluorides as heterogeneous catalysts, Prog. Solid State Chem., 26, 97 (1998). [25] T. Krahl and E. Kemnitz, The very strong solid Lewis acids aluminium chlorofluoride (ACF) and bromofluoride (ABF)-Synthesis, structure, and Lewis acidity, J. Fluor. Chem., 127, 663–678 (2006). [26] J. A. Wilkinson, Recent advances in the selective formation of the carbon-fluorine bond, Chem. Rev., 92, 505–519 (1992). [27] C. G. Krespan and V. A. Petrov, The chemistry of highly fluorinated carbocations, Chem. Rev., 96, 3269–3301 (1996). [28] M. J. Molina, F. S. Rowland, Stratospheric sink for chlorofluoromethanes – chlorine atomic catalyzed destruction of ozone, Nature., 249, 810–812 (1974). [29] E. Kemnitz and J. Winfield, in Advanced Inorganic Fluorides: Synthesis, Characterization and Applications, A. Tressaud (Ed.), pp. 367–401, Elsevier Science S.A., Amsterdam, 2000. [30] A. Hess and E. Kemnitz, Characterization of catalytically active sites on aluminum oxides, hydroxyfluorides, and fluorides in correlation with their catalytic behavior, J. Catalysis, 149, 449–457 (1994). [31] N. Herron and W. E. Farneth, The design and synthesis of heterogeneous catalyst systems, Adv. Mater., 12, 959–968 (1996). [32] J. L. Delattre, P. J. Chupas, C. P. Grey, A. M. Stacy, Plasma-fluorination synthesis of high surface area aluminum trifluoride from a zeolite precursor, J. Am. Chem. Soc., 123, 5364–5365 (2001). [33] E. Kemnitz, U. Groß, S. Rudiger, C. S. Shekar, Amorphous metal fluorides with extraordinary high surface areas, Angew. Chem. Int. Ed., 42, 4251 (2003). [34] EKLY Hajime, J. L. Delattre and A. M. Stacy, Temperature-dependent halogen-exchange activity studies of zeolite-derived aluminum trifluoride, Chem Mater., 19, 894–902 (2007). [35] A. Wander, B. G. Searle, C. L. Bailey, and N. M. Harrison, Composition and structure of the alpha AlF3 (0001) surface, J. Phys. Chem. B., 109, 22935 (2005). [36] A. Wander, C. L. Bailey, B. G. Searle, S. Mukhopadhyay and N. M. Harrison, Identification of possible Lewis acid sites on the beta-AlF3(100) surface: an ab initio total energy study, Phys. Chem. Chem. Phys., 7, 3989–3993 (2005). [37] A. Wander, C. L. Bailey, S. Mukhopadhyay, B. G. Searle and N. M. Harrison, Ab initio studies of aluminium fluoride surfaces, J. Mater. Chem., 16, 1906–1910 (2006). [38] S. Chaudhuri, P. Chupas, B. J. Morgan, P. Madden and C. P. Grey. An atomistic MD simulation and pair-distribution-function study of disorder and reactivity of alpha-AlF3 nanoparticles, Phys. Chem. Chem. Phys., 8, 5045–5055 (2006). [39] L. Francke E. Durand, A. Demourgues, A. Vimont, M. Daturi and A. Tressaud. Synthesis and characterization of Al3þ, Cr3þ, Fe3þ and Ga3þ hydroxyfluorides: correlations between structural features, thermal stability and acidic properties, J. Mater. Chem., 13, 2330–2340 (2003). [40] M. Moreno, A. Rosas, J. Alcaraz, M. Hernandez, S. Toppi, P. Da Costa, Identification of the active acid sites of fluorinated alumina catalysts dedicated to n-butene/isobutane alkylation, App. Catalysis A: General, 251, 369 (2003). [41] D. Dambournet, A. Demourgues, C. Martineau, S. Pechev, J. Lhoste, J. Majimel, A. Vimont, J.-C. Lavalley, C. Legein, J-Y Buzare´, F. Fayon and A. Tressaud, Nano-structured aluminium hydroxyfluorides derived from -AlF3, Chem. Mater., 20, 1459 (2008).

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[42] D. Dambournet, A. Demourgues, C. Martineau, E. Durand, J. Majimel, A. Vimont, H. Leclerc, J.-C. Lavalley, M. Daturi, C. Legein, J.-Y. Buzare´, F. Fayon and A. Tressaud, Structural investigations and acidic properties of high surface area pyrochlore aluminium hydroxyfluoride, J. Mater. Chem., 18, 2421 (2008). [43] D. Dambournet, A. Demourgues, C. Martineau, E. Durand, J. Majimel, C. Legein, J.-Y. Buzare´, A. Vimont, H. Leclerc and A. Tressaud, Microwave synthesis of an aluminum fluoride hydrate with cationic vacancies: structure, thermal stability and acidic properties, Chem. Mater., 20, 7095 (2008). [44] K. J. Rao, K. Mahesh, S. Kumar, A strategic approach for preparation of oxide nanomaterials, Bull. Mater. Sci., 28, 19 (2005). [45] I. V. Kubrakova, R. Khamizov, Fast determination of reaction kinetic parameters with the use of microwave heating. Kinetics of decomposition of organic substances with nitric acid, Russ. Chem. Bull., 54, 1413–1417 (2005). [46] D. S. Jacob, L. Bitton, J. Grinblat, I. Felner, Y. Koltypin and A. Gedanken, Are ionic liquids really a boon for the synthesis of inorganic materials? A general method for the fabrication of nanosized metal fluorides, Chem. Mater., 18, 3162–3168 (2006). [47] D. Dambournet, G. Eltanamy, A. Vimont, J. C. Lavalley, J. M. Goupil, A. Demourgues, E. Durand, J. Majimel, S. Rudiger, E. Kemnitz, J. M. Winfield and A. Tressaud, Coupling sol-gel synthesis and microwave-assisted techniques: A new route from amorphous to crystalline high-surface-area aluminium fluoride, Chem. Eur. J., 14, 6205–6212 (2008). [48] S. Rudiger, G. Eltanany, U. Groß and E. Kemnitz, Real sol-gel synthesis of catalytically active aluminium fluoride, J. Sol-Gel Sci. Techn., 41, 299 (2007). [49] S. Ruediger, U. Groß, M. Feist, H. A. Prescott, S. Chandra Shekar, S. I. Troyanov and E. Kemnitz, Non-aqueous synthesis of high surface area aluminium fluoride – a mechanistic investigation, J. Mater. Chem., 15, 588 (2005). [50] J. Rodriguez-Carvajal and T. Roisnel, Line Broadening Analysis Using Fullprof: Determination of Microstructural Properties, EPDIC8 Uppsala. 2002. [51] P. Thompson, D. E. Cox and J. B. Hastings, Rietveld refinement of Debye-Scherrer synchrotron X-ray data from Al2O3, J. Appl. Crystallogr., 20, 79 (1987). [52] A. Wander, C. L. Bailey, S. Mukhopadhyay, B. G. Searle, and N. M. Harrison, Steps, Microfacets, and Crystal Morphology: An ab Initio Study of -AlF3 Surfaces, J. Phys. Chem. C. 112, 6515–6519 (2008). [53] C. L. Bailey, A. Wander, S. Mukhopadhyay, B. G. Searle, and N. M. Harrison, Characterization of Lewis acid sites on the (100) surface of beta-AlF3: Ab initio calculations of NH3 adsorption, J Chem Phys., 22, 224703 (2008). [54] C.L. Bailey, A. Wander, S. Mukhopadhyay, B.G. Searle and N. M. Harrison, Adsorption of HF and HCl on the beta-AlF3 (100) surface, Phys. Chem. Chem. Phys., 10, 2918–2924 (2008). (See also Chapter 6 in this book.). [55] P. J. Chupas, D. R. Corbin, V. N. M. Rao, J. C. Hanson and C. P. Grey, A combined solid-state NMR and diffraction study of the structures and acidity of fluorinated aluminas: implications for catalysis, J. Phys. Chem. B, 107, 8327–8336 (2003). [56] P. J. Chupas, M. F. Ciraolo, J. C. Hanson and C. P. Grey. In situ X-ray diffraction and solid-state NMR study of the fluorination of g-Al2O3 with HCF2Cl, J. Am. Chem. Soc., 123, 1694–1702 (2001). [57] O. Bose, E. Kemnitz, A. Lippitz and W. E. S. Unger, C 1 s and Au 4 f(7/2) referenced XPS binding energy data obtained with different aluminium oxides, -hydroxides and -fluorides, Fres. J. of Ana. Chem., 358, 175–179 (1997). [58] O. Boese, W. E. S. Unger, E. Kemnitz and S. L. M. Schroeder, Active sites on an oxide catalyst for F/Cl-exchange reactions: X-ray spectroscopy of fluorinated -Al2O3, Phys. Chem. Chem. Phys., 4, 2824–2832 (2002). [59] D. Dambournet, H. Leclerc, A. Vimont, J. C. Lavalley, M. Nickkho-Amiry, M. Daturi and J. M. Winfield, The use of multiple probe molecules for the study of the acid-base properties of aluminium hydroxyfluoride having the hexagonal tungsten bronze structure: FTIR and [36Cl] radiotracer studies, Phys. Chem. Chem. Phys., 11, 1369–1379 (2009). [60] R. Chandross, The structure of a new phase of aluminum trifluoride monohydrate, Acta Cryst. 17, 1477 (1964). (It should be noted that the synthesis of a cubic form of an aluminium fluoride monohydrate proposed by Chandross could not be satisfactorily reproduced.).

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[61] E. Kemnitz, U. Groß, St. Ru¨diger, G. Scholz, D. Heidemann, S.I. Troyanov, I.V. Morosov and M.-H. Leme´e-Cailleauc, Comparative structural investigation of aluminium fluoride solvates, Solid State Sci., 12, 1443–1452 (2006). [62] L. F. Scatena, M. G. Brown and G. L. Richmond, Water at hydrophobic surfaces: weak hydrogen bonding and strong orientation effects, Science, 292, 908–910 (2001). [63] K. S. Vorres and F. B. Dutton, The fluorides of titanium: X-ray powder data and some other observations, J. Am. Chem. Soc., 77, 2019 (1955). [64] K. Vorres and J. Donohue, The structure of titanium oxydifluoride, Acta Cryst., 8, 25–26 (1955). [65] A. Demourgues, N. Penin, E. Durand, F. Weill, D. Dambournet, N. Viadere and A. Tressaud, New titanium hydroxyfluoride Ti0.75(OH)1.5F1.5 as UV absorber, Chem. Mater., 21, 1275–1283 (2009). [66] Y. Xie, S. Yin, H. Yamane, T. Hashimoto, H. Machida and T. Sato, Microwave assisted solvothermal synthesis of a new compound, pyrochlore-type Sn1.24Ti1.94O3 66(OH)1.50F1.42, Chem. Mater., 20, 493–495. (2008). [67] G. Ferey, Hybrid porous solids: past, present, future, Chem Soc Rev., 37 191–214 (2008). [68] H. Jhung, J.-H. Lee, J. W. Yoon, C. Serre, G. Fe´rey and J.-S. Chang. Microwave synthesis of chromium terephthalate MIL-101 and its benzene sorption ability, Adv. Mater., 19, 121–124 (2007). [69] A. Demourgues, L. Sronek and N. Penin, New Nanostructured Flurocompounds as UV Absorbers, this book, Chapter 8, pp. 229–273, John Wiley & Sons (2010). [70] N. Penin, N. Viadere, D. Dambournet, A. Tressaud, and A. Demourgues, Tuned optical band gap for titanium-based oxy(hydroxyl)fluorides, Proceedings MRS Fall Meeting, Boston, USA, November 2005. [71] N. Penin, N. Viadere, D. Dambournet, A. Demourgues, and A. Tressaud, Synthesis and characterization of Ti-based oxy-hydroxy-fluorides, Proceedings 18th International Symposium on Fluorine Chemistry, Bremen, Germany, July 2006.

3 High Surface Area Metal Fluorides as Catalysts Erhard Kemnitz and Stephan Ru¨diger Humboldt-Universita¨t zu Berlin, Institut fu¨r Chemie, Brook – Taylor – Str. 2, D – 12489 Berlin, Germany

3.1

Introduction

Metal fluorides prepared via the sol-gel fluorination route are nanoscopic and consequently have a high surface area. Metal ions at the surface of a particle cannot have a coordination that is as symmetric as that of ions inside the particle. They are coordinatively undersaturated. That effect together with the strong electron-withdrawing power of fluorine brings about the inherent Lewis acidity of the respective metal ions. Thus, high surface-area aluminium fluoride, HS-AlF3, is an exceptionally strong Lewis acid, as shown in Chapter 1; its acidity corresponds to its position in the Lewis acidity scale calculated for isolated molecules, where AuF5 and SbF5 are the strongest Lewis acids followed by AlCl3 > AlCl2F > AlClF2 > AlF3 although all these aluminium halides exhibit nearly the same acidity [1]. It is interesting to note that magnesium fluoride, typically regarded as neutral compound, also shows distinct Lewis acidity, although not as strong by far as that of HS-AlF3, when prepared via the sol-gel fluorination route [2]. The fact that the acidity of HS-MgF2 diminishes upon heating, as follows from the respective ammonia TPD investigations (Figure 3.1) proves the dependence of the Lewis acidity on the low structural order of the material. Chapter 1 describes how, on one hand, the Lewis acidity of high surface metal fluorides can be adjusted

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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and how, on the other hand, they can be modified to alter or add other properties than Lewis acidity. Most of these variations have been tested for or have even been used with the aim of improving catalytic properties for selected reactions with often surprisingly good results. These exciting catalytic properties of nanoscopic metal fluorides obtained via fluorolytic sol gel synthesis will be presented in the following. Based on their acidity, often expressed in terms of catalytic activity, the catalyst in the following will be qualitatively differentiated as strong, middle strong, weak, etc. Unfortunately, there is no objective quantitative measure for the strength of acidic surface sites that could be used to rank solid acids as it is well established for acids in aqueous systems by the KS values. On the other hand, FTIR-spectra of probe molecules, e.g. pyridine or CO, adsorbed on acidic surface sites can be used as a measure based on the shift of the absorption peak. The most commonly used probe molecule for routine measurements is pyridine. However, since this is a strong base, its shift of the absorption peak in the FTIR spectrum is too small to give a serious quantification of the acidic strength of a surface site. It is nevertheless an excellent probe molecule to differentiate between the kind of acidity, 32

pretreated at 773 K

30

pretreated at 573 K pretreated at 393 K

28 26 24 22

Intensity/a.u.

20 18 16 14 12 10 8 6 4 2 0 –2 300

400

500

600

700

800

800

Temperature/K Figure 3.1 NH3-TPD of HS-MgF2 preheated at different temperatures. (Reprinted with permission from [2] Copyright (2006) Elsevier Ltd.)

High Surface Area Metal Fluorides as Catalysts

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that is Brønsted or Lewis, which will be detailed at the appropriate places of this chapter. Carbon monoxide, on the other hand, is a very sensitive probe molecule in terms of peak shift depending on the strength of the acidic surface site. However, since CO is a very weak basic molecule, it is less sensitive regarding weak acidic sites. Alternatively, temperature programmed desorption (TPD) of ammonia gives a comprehensive figure about the distribution of acid strength of surface sites. The principle of this method is, in brief, as follows. The respective solid sample is deposited in a suitable flow reactor, which can be controlled and heated and which is connected with an FTIR-cell. On the precalcined sample (usually between 523 and 573 K) the first dry ammonia is adsorbed at e.g. 373 K. After saturation with ammonia the sample is then flashed by an inert gas stream passing through the connected FTIR cell until ammonia that is no longer desorbed can be detected. Now the reactor will be subjected to a programmed temperature raise and the desorbed ammonia is followed as a function of temperature. Thus a concentration versus temperature profile will be obtained (see for example Figure 3.1) that gives a reflection of the acid site distribution in which the strength of sites is indicated by the temperature (the higher the temperature the stronger the sites) and the concentration of sites of a certain strength is proportional to the intensity of the IR signal. By absorbing the desorbed ammonia by an aqueous acid, the overall amount of desorbed ammonia can be exactly determined that corresponds to the square below the desorption graph. Based on this one may go to determine different ranges of acidic strengths. This might be acceptable as long as different solids that are similar in nature (e.g. metal fluorides) are compared and as long as the same equipment is used. However, since these profiles depend on several parameters (surface morphology, design of the desorption cell, heating rate, detection system etc.) such data would not really allow any serious comparison of different samples measured in different labs. Hence, although both of the techniques mentioned above are very often used – and are also used by the authors – it seems less useful to apply real quantification because of the problems briefly mentioned. If not otherwise stated the different catalytic parameters used in this chapter have the following meaning: Conversion defines the part of converted compound A (n0AnA) in relation to the starting concentration (n0A); X ¼ (n0AnA)/n0A Selectivity defines the part of a desired product P among all the products formed: SP ¼ ðnP veduct Þ=nP ðn0educt neduct Þ Yield defines the part of the desired product in the reaction mixture, YP ¼ SPXA

3.2

High Surface Area Aluminium Fluoride as Catalyst

The very high Lewis acidity of HS-AlF3 causes it to react easily with most electron pair donating solvents such as alcohols or aqueous ones and many reactants whereby the acidic sites of HS-AlF3 become blocked. Therefore, if its high Lewis acidity is of interest for catalytic reactions it can be used at best in fluoro-organic reactions although a few others have also been reported. Seven reactions of fluoro-organic compounds have been successfully catalysed by HS-AlF3 (Equations (3.1) to (3.7)) [3], the use of the reactions in

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Equations (3.1) to (3.4) as test reactions to assess the Lewis acidity of the catalyst is already described in Chapter 1. 5 CCl2 F2 ! CCl3 F þ 3 CCl3 F þ CCl4

(3:1)

3 CHClF2 ! CHCl3 þ 2 CHF3

(3:2)

CBrF2 CBrFCF3 ! CF3 CBr2 CF3

(3:3)

CCl2 FCClF2 ! CCl3 CF3

(3:4)

F3 C-CHF-CHF2 ! F3 C-CF¼CHF þ HF

(3:5)

F3 C-CH2 -CF3 ! F3 C-CH¼CF2 þ HF

(3:6)

F3 C-CH2 F ! F2 C¼CHF þ HF

(3:7)

The dismutation reactions (3.1) and (3.2) proceed only under the catalytic action of a Lewis acid albeit a moderate Lewis acidity is sufficient. Thus, HS-MgF2 is reported to catalyse the dismutation reaction (Equation (3.2)) in a flow reactor with 3% conversion at 573 K and 60% conversion at 623 K whereas common MgF2 is not at all active [2]. With the much stronger Lewis acid HS-AlF3 reactions (Equation (3.1)) and (Equation (3.2)) proceed with almost 100% conversion even at room temperature [4]. The isomerization reaction (Equation (3.3)) proceeds only under the catalytic action of the strongest Lewis acids known, SbF5 (at elevated temperatures, 383 K), aluminium chlorofluoride (ACF), and HS-AlF3 [5]. As with ACF, with HS-AlF3 almost 100% conversion was observed at room temperature in a batch reaction. The isomerization reaction (Equation (3.4)) also proceeds very effectively in the presence of HS-AlF3. In a detailed study comparing different Lewis acidic Al-F and Al-Cl based catalysts, which is described in Chapter 1, HS-AlF3 proved to be almost equal to ACF and superior to the AlCl3 and -AlF3, which are both not active at room temperature. However, -AlF3 is well known as a good and selective catalyst for this reaction at temperatures above 553 K [6–8]. AlCl3 however, became very active after being converted into ACF by a chemical reaction with CCl2FCClF2 [5]. Hence, the latter experiments proved that solid HS-AlF3 is a better Lewis-acid catalyst than solid AlCl3. The well known excellent Lewis acidity of AlCl3, – and based on this its high catalytic activity in many organic reactions for which in contrast the currently known AlF3 phases are totally inactive – is just a result of the better solubility of AlCl3 in comparison to the practically insoluble AlF3. Thus, the catalytic difference between these two crystalline phases is exclusively caused by their different solubility in organic solvents. Although HF behaves towards HS-AlF3 as base adsorbed at and blocking the strongest Lewis acid sites [7], HS-AlF3 has also proved to be an effective catalyst for dehydrofluorination reactions (Equations (3.5) to (3.7) [3]. It is known that adsorbed HF is almost completely released at above 573 K as can be seen in Figure 3.2 [7]. Therefore, such reactions have to be carried out at temperatures distinctly higher than those needed for the reactions in Equations (3.1) to (3.4). At 623 K 1,1,1,2,3,3-hexafluoropropane is dehydrofluorinated with 95% conversion to the 1,1,1,2,3-pentafluoropropene with 88% selectivity for the cis-isomer and 12 to the trans-isomer (Equation (3.5)). No formation of any 1,1,1,3,3-pentafluoropropene was

High Surface Area Metal Fluorides as Catalysts DTG /(% min–1) DTA /μV Ion Current ·10–10/A

TG (%) 100.0 TG

10

818

95.0 DTG

0

90.0 85.0

73

DTA ↑ exo

–14.43 –10

m19(×40)

–20

80.0

–30

806

75.0

–40

70.0

m18

65.0

m20(×40)

–60

473

573

673 773 Temperature/K

8.0 7.0

–0.10 6.0 –0.20

5.0

–0.30

4.0

–0.40

3.0

–0.50

2.0

–0.60

1.0

–0.70

0

–50

60.0 373

0

873

Figure 3.2 TA/MS curves of HS-AlF3 loaded previously with HF, and IC curves for m/z 18 (H2Oþ), m/z 19 (Fþ), and m/z 20 (HFþ). The identical curve shapes for m/z 19 and 20 can be seen as proof that they originate from the same molecule, i.e. HF, whereas m/z 18 (H2Oþ) is different. (Reprinted with permission from [7] Copyright (2005) Royal Society of Chemistry.)

observed under these conditions. Under similar conditions at 673 K 1,1,1,3,3,3-hexafluoropropane reacts to 1,1,1,3,3-pentafluoropropene with 33.5% conversion and 66% selectivity towards 1,1,1,3,3-pentafluoropropene (Equation (3.6)). With decreasing temperature in both reactions the conversion decreases but the selectivity increases. Dehydrofluorination of 1,1,1,2-tetrafluoroethane to trifluoroethene (Equation (3.7)) also proceeds at 623 K with 20% conversion and 78% selectivity toward the desired product. The reactions in Equations (3.6) and (3.7) necessitate an attack on the respective CF3 group, recognized to be very stable. The activation of these groups within the progression of the reactions gives evidence for the unusually high Lewis acidity of the catalyst although this was to a minor extent also observed at fluorinated crystalline aluminium fluorides and fluorinated aluminas [8]. Other reactions reportedly catalysed by HS-AlF3 are shown in Equations (3.8) to (3.10).

(3.8)

+ H H

H

H

+

H H

(3.9)

ðCH3 Þ 3 CCl ! CH2 ¼CðCH3 Þ2 þ HCl; CH2 ¼ CðCH3 Þ2 ! CH3 CH ¼ CHCH3 (3:10) Isomerization reactions (3.8) and (3.9) have been carried out at 573 K or 473 K, respectively, in a flow reactor, whereby in reaction (3.8) 83% conversion and in reaction (3.9) 47% conversion was observed [3]. Reaction (3.10), dehydrochlorination of tBuCl

74

Functionalized Inorganic Fluorides

followed by isomerization of the 2-methylpropene to butene-(2), was observed at room temperature in course of radiotracer experiments and is described in Chapter 4. The before mentioned isomerisation reaction of 2-methylpropene to butene-(2) was observed for the first time under such circumstances indicating the very strong Lewis acidity of HS-AlF3. A solid catalyst designed for use in a continuous-flow reactor has to meet some physical requirements, such as mechanical stability and low flow resistance, requirements that are not a priori fulfilled by HS-AlF3, which is a fine powder. Fortunately, there is a convenient way to support HS-AlF3 on g-alumina (see Chapter 1), and the supported catalyst showed similar catalytic activity in all the above reactions [10].

3.3

Host-Guest Metal Fluoride Systems

According to the Tanabe model [11], originally developed for metal oxide host-guest systems but later adapted for metal fluoride host-guest systems [12], the kind of acidity (Lewis or Brønsted) can be predicted based on the anion coordination of the host and guest metal ions, respectively. This is a pure geometric model but allows with up to 90% accuracy to predict whether Brønsted or Lewis acidity will be generated by creating solid solutions. However, it does not give any prediction of the strength of acidity. In brief, the following assumptions are made according to this approach applied for a solid solution two different metal fluorides (MFx and M’Fy). (i) The coordination number of both of the metal cations in MFx and M’Fy are maintained even when mixed. (ii) The coordination number of the fluoride ion of the major (host) metal fluoride component is retained for all the fluorides in a ternary metal fluoride system. Suppose, for example, that CrF3 is introduced as guest in the MgF2-host system. The coordination numbers of both metals are 6 and remain even in the guest-host system, whereas the coordination number of all the fluoride ions is 3. Keep in mind that the CN of F in CrF3 is just 2! If we now count the electrostatic bonding power, we have the following situation: the positive charge of Mg is 2þ and will be distributed to six surrounding (coordinating) F-atoms, i.e. þ2/6 ¼ þ1/3 charge is directed to each F-atom. On the other hand, the charge of F is distributed towards three coordination Mg atoms, e.g. –1/3. This means that the charges are perfectly balanced between Mg and F as expected. However, the situation is different in case of the host cation Cr3þ: since Cr3þ is coordinated by 6 F atoms, the charge is distributed by þ3/6 ¼ þ1/2 to each surrounding F atom. However, since all the F atoms exhibit the same coordination number 3 the charge coming from the F atom is just –1/3. For each Cr3þ cation the overall charge balance can be calculated by 6x (þ1/2 – 1/3) ¼ þ1. The conclusion from this is that by introducing CrF3 into a MgF2 host a local positive charge at the chromium site is generated acting as a Lewis acidic site. In an opposite way (introduction of MgF2 into a CrF3 host), negative charges will be generated which becomes compensated by attraction of protons thus creating Brønsted sites (see details [12]). Mixed metal fluoride systems where the cations of the minor component can principally substitute cations of the major component in the lattice therefore open an easy way to

High Surface Area Metal Fluorides as Catalysts

75

Table 3.1 Benzoylation/acetylation of anisole, i-propylation of benzene and dismutation of CCl2F2 over host/guest catalysts (Data taken from [13] by permission of Elsevier Publishers) Catalyst

Fe/MgF2 Ga/MgF2 In/MgF2 V/MgF2 a b c

Benzoylation with Ph-COCl

Acetylation with AcOAc

Xb(%)

Yc(%) 4-MBPh

X (%)

100 100 100 0

69 77 79 0

45 15 nda 9

Y (%) 4-MAPh 34 6.5 – 1

i-propylation with CH3CHClCH3

Dismutation of CCl2F2

X Y (%)-cumene (%)

Temp. (K)

X (%)

100 nd nd nd

623 623 623 623/573 523/473

25 30 5 85/56/ 28/10

84 – – –

nd: not determined. X: conversion. Y: yield.

create metal fluoride-based solid catalysts with adjusted Lewis acidity (see Chapter 1). The necessity for adjusted Lewis acidity follows from a comparison of HS-AlF3 and -AlF3 as solid catalysts. Although HS-AlF3, as very strong Lewis acid, should be an ideal catalyst for all Lewis acid-catalysed reactions, it cannot be used in many of these reactions because it becomes almost immediately deactivated by often irreversible adsorption of educts and/ or necessary solvents thus blocking the Lewis acid sites. -AlF3 on the other hand is not a strong enough Lewis acid for many reactions. Many MgF2/MF3 systems are reported with HS-MgF2 being the major component, the host, and the guest metal fluoride amounts to up to 30%. Such systems, which should be rather considered as solid solutions and which might be doped systems, have been reported with Fe3þ, Ga3þ, V3þ, In3þ, and Cr3þ as guest cations [13–15]. For MgF2 the precursor was always an alkoxide, whereas fortunately, for the guest component a broader range of compounds could be used as precursor. It was found that surface area and pore size as well as Lewis acidity and catalytic activity depend on the nature of the guest cation and its concentration, as expected, and also partly on the chemical nature of the compound that was used to introduce the guest metal during the sol-gel synthesis. The effect of the respective guest cation follows from Table 3.1. All these catalysts were prepared from Mg(OMe)2, MCl3, and In(OiPr)3 (85:15), respectively [13]. The catalysts showed very different activity for dismutation of CCl2F2 (Equation (3.1)) despite only moderate variations in surface area (100155 m2/g), and the best one, V/MgF2, which is clearly inferior to HS-AlF3, was on the other hand not at all active in the benzoylation reaction. Obviously, the Lewis acidity of V/MgF2 is that high, as proved by its dismutation activity, that at the acid sites of the solid components of the benzoylation reaction adsorb irreversibly thereby blocking the catalyst as known for HS-AlF3. It is, however, worth noting that the simple solid Fe3þ/MgF2 catalyst was rather active for Friedel-Crafts benzoylation (Equation (3.11)) and acylation as well as alkylation (Equation (3.12)). CH3 OC6 H5 þ C6 H5 -CðOÞCl ! CH3 OC6 H4 -CðOÞ-C6 H5 þHCl cat

C6 H6 þ C3 H7 Cl ! C6 H5 -C3 H7 þHCl

(3:11) (3:12)

76

Functionalized Inorganic Fluorides

In a more detailed investigation of the Fe3þ/MgF2 catalyst system (Fe:Mg ¼ 15:85) three types of Fe3þ precursor, i.e. FeCl3, Fe2(SO4)3 and Fe(OCH3)3, and two ways of postfluorination, i.e. with CCl2F2 or with HF, were employed [14]. The surface areas of the HF fluorinated catalysts (174–298 m2/g) were markedly higher than of the CCl2F2 fluorinated ones (75–104 m2/g). The latter showed, however, a somewhat higher catalytic activity for CCl2F2 dismutation (Equation (3.1)) and especially anisole benzoylation (Equation (3.11)) with 97–100% conversion and 60–82% selectivity whereas the HF fluorinated catalysts gave 41–89% conversion and 42–62% selectivity. An Fe(OCH3)3 based and HF postfluorinated catalyst was repeatedly successfully used in benzoylation, giving no evidence of leaching. Two catalysts of the CCl2F2 postfluorination group, based on FeCl3 (named MgFeClR) and Fe2(SO4)3 (named

B 90

553 K

593 K

633 K

80 70 60 50 40 30

Relative content (%)

20 10 0 conv

R123

R124

R125

others

conv

R123

R124

R125

others

A 70 60 50 40 30 20 10 0

Figure 3.3 Hydrofluorination of C2Cl4 over MgFeClR (lower graph, obtained from FeCl3 and postfluorinated by CCl2F2) or over MgFeSR (upper graph, obtained from Fe2(SO4)3 and postfluorinated by CCl2F2). (Reprinted with permission from [14] Copyright (2004) Elsevier Ltd.)

High Surface Area Metal Fluorides as Catalysts

Conversion of CFC12, %

90 80

CrClMg15 CrOMg8

CrAcMg15 CrOMg25

77

CrOMg15

70 60 50 40 30 20 10 0 350

400

450

500 550 Temperature, K

600

650

Figure 3.4 Dismutation of CCl2F2 over CrF3/MgF2 catalysts at tR ¼ 2 s; CCl2F2:N2 ¼ 1:5. (Reproduced from [15] by permission of Elsevier.)

MgFeSR), respectively, have also been tested for hydrofluorination of C2Cl4 (Equation (3.13)). CCl2 CCl2 þðx þ 1ÞHF ! C2 HCl4 x Fxþ1 þxHCl

(3:13)

The technically useful reaction (3.13) cannot be catalysed by HS-AlF3 because of irreversible adsorption of HF. Therefore, it is a very interesting result that with the Fe3þ/ MgF2 catalysts, which are highly Lewis acidic although not as high as HS-AlF3, the hydrofluorination proceeded well (Figure 3.3) [14]. There is a similarly detailed study of the Cr3þ/MgF2 guest/host system scrutinizing physicochemical characteristics and catalytic properties of both, differently prepared and composed solids [15]. The catalysts were prepared from Mg(OMe)2 and CrCl3 (Mg:Cr ¼ 85:15; named CrClMg15), Cr3Ac7(OH)2 (85:15; CrAcMg15)) and CrO3, prior to use reduced by reaction with MeOH, (82:8, 85:15 and 75:25; CrOMg8, CrOMg15 and CrOMg25). CCl2F2 was employed as postfluorination agent. The XRD patterns of all catalysts show very broad reflexes only for MgF2 and give no hint for other phases. Their surface area ranged from 65 m2/g (CrAcMg15) over 99 m2/g (CrOMg25) and 106 m2/g (CrClMg15) to 167 m2/g (CrOMg8) and 175 m2/g (CrOMg15). CrOMg15 had the highest surface area and was the most active catalyst for CCl2F2 dismutation (Figure 3.4), indicating the highest Lewis acidity of all the preparations. Comparing the CrOMg catalysts of different molar Cr:Mg ratio it becomes evident that about 15 mol% guest cations is the optimum composition. This holds also for the dismutation of CHClFCF3 according to Equation (3.14) the results of which are shown in Figure 3.5. 2CHClFCF3 ! CHF2 CF3 þCHCl2 CF3

(3:14)

Obviously, guest cations up to a concentration of about 15 mol% can readily become incorporated into the lattice of the host, provided the two cations are of comparable size, enhancing the Lewis acidity and the derived catalytic activity of the solid solution compared to the neat host system.

78

Functionalized Inorganic Fluorides 100

CHCl2CF3

CHClFCF3

Relative concentration (%)

90 80 70 60 50 40 30 20 10 0

CrClMg15 CrAcMg15 CrOMg15

CrOMg8

CrOMg25

Figure 3.5 Dismutation of CHClFCF3 over Cr/MgF2 catalysts at 573 K; tR ¼ 2 s; CHClFCF3:N2 ¼ 2:1.5. (According to Equation (3.14), the amounts of CHF2CF3 have to correspond with the respective amounts of CHCl2CF3. However, the experimental values for CHF2CF3 are smaller due to losses because of its higher volatility. They are therefore omitted.) (Reprinted with permission from [15] Copyright (2005) Elsevier Ltd.)

3.4

Hydroxy(oxo)fluorides as Bi-acidic Catalysts

Depending on the nature of the metals the M-OH units created in a metal fluoride system may either react as a base or as a (Brønsted) acid. Thus, replacement of F by OH, or vice versa, should result in an alteration of the basicity of M-OH groups, e.g. in MgF2x(OH)x derivatives as compared to Mg(OH)2 and in the same way of the (Brønsted) acidity, e.g. in AlF3x(OH)x phases as compared to Al(OH)3. Thus, some of the fluoride ions of a metal fluoride can be replaced by hydroxyl and/or oxide groups in the course of its preparation by introducing defined amounts of water at different steps of the synthesis, modifying the properties. A number of magnesium fluoride-derived materials have been prepared starting from Mg(OMe)2 and understoichiometric amounts of HF, which immediately form a sol to which twice the theoretically needed amount of water was added. One preparation started from a stoichiometric Mg:HF ratio with supplementary water added to the sol. The final gels were dried and calcined at 350 C [16]. The preparations, named MOF-0, MOF-0.4, MOF-1.2, MOF-1.6, and MOF-2.0 according to relative amount of HF to Mg used, had analytically about 80% of the nominal F content, obviously due to incomplete fluorination in the gel phase, as is known for HS-MFx syntheses. Despite their designations the hydroxyl groups introduced are almost completely dehydrated at the calcination temperature, creating oxidic O atoms. The materials showed graded physico-chemical properties in terms of surface area, ranging from 104 m2/g (MgF2) to 387 m2/g (Mg(OH)1.2F0.8), and in terms of XRD pattern, FTIR spectra, 19F MAS NMR spectra as well as Mg 1s, F 1s and O 1s binding energies [16]. The results of Michael addition experiments are very interesting, i.e. addition of

High Surface Area Metal Fluorides as Catalysts

79

2-methyl-cyclohexane-1,3-dione to methyl vinyl ketone, employing the synthesized magnesium hydroxide fluorides as catalyst compared to MgO and crystalline MgF2 catalysts, shown in Figure 3.6. MgF2 was not at all active (so it is not included in Figure 3.6), whereas MgO was very active giving 70% yield in less than 1 h. However, after 1 h the yield started to decrease to about 20% in 24 h caused, by consecutive reactions of the primary addition product. Likewise, with commercial Mg(OH)2 a maximum yield of 70% is reached after 4 h, followed also by a decline to 45% in 24 h. The F-containing catalysts showed a more differentiated activity. The catalyst with the lowest F content, h-Mg(OH)1.6F0.4, was the

A 100 MOF-0 MOF-0.4 MOF-1.2 MOF-1.6 MOF-2.0 Mg(OH)2

Yield of Michael adduct (%)

90 80 70 60 50 40 30 20 10 0 0

2

4

6

8

20

22

24

Time/h

Selectivity of Michael adduct (%)

B

100 MOF-0 MOF-0.4 MOF-1.2 MOF-1.6 MOF-2.0 Mg(OH)2

90 80 70 60 50 40 30 20 10 0 0

2

4

6

8

20

22

24

Time/h

Figure 3.6 Michael addition of 2-methyl-cyclohexane-1,3-dione to methyl vinyl ketone catalysed by Mg(OH,F)2 catalysts. Yield (above) and selectivity (below) versus reaction time. (Reprinted with permission from [16] Copyright (2005) Elsevier Ltd.)

80

Functionalized Inorganic Fluorides

most active of all, with about 90% yield after 2 h, which, however, decreased rapidly as in case of MgO, indicating that there are basic sites at the catalyst that are too strong, promoting consecutive reactions. The two catalysts with the highest F content, hMg(OH)0.4F1.6 and h-MgF2, showed a totally different activity – they were less active but gave the highest yield after 24 h because of absence of consecutive reactions. Accordingly, they were very selective (Figure 3.6). The most active catalyst h-Mg(OH)0.4F1.6 was also used in three other Michael addition reactions, i.e. addition of 2-acetylcyclopentanone, 2-acetylcyclohexanone, and 2-methoxycarbonyl-cyclopentanone to methyl vinyl ketone. These reactions proceeded also with high (75–90%) yields and 100% selectivity [16]. The great dependence of activity on the F content of the samples was explained as follows. Both, F– and O2– are by definition Lewis bases but pure MgF2 did not give any reaction at all, meaning the F-sites do not act as Lewis basic sites. In contrast the magnesium oxofluorides carrying O2 are very active. Hence, the explanation is that O2 sites at the solid surface are excellent proton acceptors forming Brønsted basicity, which is necessary for the catalytic performance but F–-ions are not at all. The F atoms, on the other hand, reduce the basicity of the O atoms due to their electron-withdrawing effect. Thus, by changing the F/O ratio the basicity of the O atoms can be tuned down to a level sufficient for catalysing the addition reaction but too low for the consecutive reactions. Another closely related synthesis route to partly hydroxylated metal fluorides starts from metal alkoxide in alcohol, which is reacted with the stoichiometric amount of hydrogen fluoride dissolved in different amounts of water thus enabling the competition between fluorolysis and hydrolyis, which to some extent is ruled by the HF to H2O ratio and results in partly hydroxylated fluorides. Thus, Mg(OCH3)2 dissolved in MeOH was reacted with stoichiometric amounts of aqueous HF as 40wt%, 57wt%, 71wt%, and 87wt% solution yielding a viscous, transparent gel, which was dried under vacuum firstly at room temperature and finally at 70 C for 5 h. The hydroxylated MgF2 products, named MgF2-40, MgF2-57, MgF2-71 and MgF2-87 and having surface areas ranging from 180 to 420 m2/g, have been employed as catalyst for (all-rac)-[]-tocopherol (vitamin E, 3) synthesis through condensation of trimethylhydroquinone (TMHQ, 1) with isophytol (IP, 2) (Equation (3.15)), a reaction known to need acid catalysis [17]. HO

HO

+ OH TMHQ, 1

OH IP, 2

C14H29

catalyst –H2O

O

C14H29

(3.15)

3

Table 3.2 shows some results of tocopherol synthesis obtained with different MgF2 based catalysts. Both crystalline MgF2 (entry 1) and HS-MgF2 prepared with very little or no water (entries 7 and 8) were not at all active, even after prolonged reaction time. As crystalline MgF2 exhibits almost no acidity its inactivity was to be expected. On the other hand, a HS-MgF2 catalyst prepared with 71% aqueous HF resulted in total conversion of isophytol and almost 100% selectivity to (all-rac)-[]-tocopherol (entry 6). The activities of the catalysts do not correspond to their respective numbers of acid centres (Table 3.2). Likewise, the 19F MAS NMR spectra do not correspond to the activity.

High Surface Area Metal Fluorides as Catalysts

81

Table 3.2 Influence of the key surface features and the reaction parameters on the catalytic performances of tocopherol synthesis Entry Catalyst 1 2 3 4 5* 6** 7 8

Number of A.C./m2

MgF2-C MgF2-40 MgF2-57 MgF2-71 MgF2-71 MgF2-71 MgF2-87 MgF2-100

n.d. 5.4 6.6 3.6 3.6 3.6 4.8 8.4

IP/catmolar Time, (min) Conversion of ratio isophytol n.d. 119 76 123 31 123 60 37

1017 1017 1017 1017 1017 1017 1017

1800 300 300 300 60 180 360 360

Yield of tocopherol (%)

0 100 100 100 100 100 0 0

0 76.3 82.6 87.0 92.6 > 99.9 0 0

Reaction conditions: 50mg of catalyst; T ¼ 373 K; TMHQ/IP ¼ 1/1; solvent: heptane/propylene carbonate¼ 50/50 (v/v). Yield is based on IP (C ¼ 100%).* 200mg of catalyst.** TMHQ/IP ¼ 2/1; A.C. ¼ acidic centres calculated from NH3-TPD and N2 sorption isotherms.

BPy

LPy

L.A.

MgF2-100 MgF2-87 MgF2-71 MgF2-57 B.A.

MgF2-40

2000

1800

1600 1400 Wavenumber [cm–1]

1200

Figure 3.7 Infrared spectra of pyridine adsorbed on HS-MgF2 preparations

The resonance signal (–198 ppm) is at almost identical position compared to that of crystalline MgF2, but with increasing water content the signal showed a slight deformation to low field, which can be assigned to increased structural disorder due to increased number of OH groups [18]. The IR spectra of adsorbed pyridine (Figure 3.7) in the region characteristic for adsorption on Brønsted acid sites and on Lewis acid sites are more informative. Thus, MgF2-100 and MgF2-87 exhibit almost only Lewis acid sites, responsible for the shoulder at 1611 cm1. This can be seen in an even more pronounced way from the intensity of the peak at ca. 1450 cm1 which decreases with increasing water content of the HF used for the synthesis. At the same time, the peak at 1545 cm1 increases. Obviously, Brønsted sites are increasingly created at the expense of Lewis sites. With increasing water content, the Lewis acid sites

82

Functionalized Inorganic Fluorides

diminish and Brønsted sites occur, the ratio of Lewis to Brønsted sites is adjusted and ‘biacidic’ material is created. Evaluating the experimental results (Table 3.2) one can see that the predominantly Brønsted acidic MgF2-40 catalyst is highly active (100% conversion) but less selective. On the contrary, the predominantly Lewis acidic MgF2-100 and MgF2-87 catalysts are not active. Obviously, the graded Brønsted acidity of the ‘bi-acidic’ MgF2-71 catalyst provides optimum conditions for the overall reaction, preventing the formation of consecutive products. The observed high regioselectivity to (all-rac)-[]-tocopherol might be due to the structural peculiarities of the catalyst as it is reported that very bulky polymerbound catalysts improve the regioselectivity of this reaction [19]. Comparing the properties of the hydroxylated magnesium fluoride catalysts, a relationship between the extent of hydroxylation and Brønsted acid/base behaviour becomes obvious. Contrary to expectations, Mg-OH bonds can be Brønsted acidic when created to a very small extent in an MgF2 host phase, obviously, due to the electron-withdrawing effect of the fluoride ions. However, the higher the degree of hydroxylation the lower the Brønsted acidity or in other terms the higher the Brønsted basicity. Very small amounts of hydroxyl groups are strong Brønsted acids as proved for tocopherol synthesis, and very high numbers of hydroxyl groups corresponding with smaller amounts of fluoride are Brønsted bases as proved for the Michael addition reactions. A simple test with aqueous acid/base indicator showed exactly the same relation. [18]. In the light of the findings with hydroxylated MgF2 catalysts, HS-AlF3 with its exceptional high Lewis acidity but small Brønsted acidity should be a poor catalyst for tocopherol synthesis. In fact, after as long as 1200 min reaction time (conditions as given in Table 3.2) only 70% IP conversion were detected, but the reaction stopped at an intermediate stage at phytylhydroquinone, PHQ. The transformation of PHQ to (all-rac)-[]-tocopherol could not be provoked by the strong Lewis and weak Brønsted acid HS-AlF3 [20]. However, with a partly hydroxylated catalyst (‘AlF3-H’) prepared from aluminium isopropylate reacted in propanol with stoichiometric amounts of HF added as 50% (w/w) solution in water and followed by drying the gel formed under vacuum (finally at 70 C for 5 h) the conversion of IP was completed within 1 h and (all-rac)-[]-tocopherol formed in very high yield. Obviously, it is not strong Lewis acid sites but moderate to medium-strong Brønsted acid sites that are needed to catalyse the overall reaction. The difference in the acidities of HSAlF3 and AlF3-H follows from the respective IR spectra of adsorbed carbon monoxide gas shown in Figure 3.8. Compared to neat HS-AlF3 AlF3-H does not have super-strong Lewis acid sites indicated by the almost complete absence of CO IR absorption bands in the region above 2200 cm1, which is indicative for very strong Lewis acid sites. The strongest absorption occurs in the region between 2160 and 2180 cm1, the intensive band at 2172 cm1 can be assigned to Brønsted acid sites, proving the dominance of Brønsted acidity. This is supported by the spectra of adsorbed lutidine in Figure 3.9 In the spectra there are two bands at 1655 and 1631 cm1 corresponding to protonated lutidine, observable only with moderate to strong Brønsted acid sites at the solid, and in case of high lutidine partial pressure (Figure 3.9a) there are additional absorption bands at 1581, 1595 and 1604 cm1 corresponding to physisorbed and weakly coordinated species, respectively. The results obtained with AlF3-based catalysts confirm the findings with hydroxylated MgF2 catalysts. They show that introducing a graded hydroxylation upon the sol-gel fluorination synthesis opens possibilities for fine-tuning the strength and balance of Lewis/Brønsted acidity and thus extends substantially the fields of application of metal fluoride-based catalysts.

High Surface Area Metal Fluorides as Catalysts

83

Adsorption

2220

2200

2160

2180

2140

Wavenumbers /cm–1

Figure 3.8 Infrared spectra of increased doses of CO adsorbed on AlF3-H at 100 K. (Reprinted with permission from [20] Copyright (2008) Wiley-VCH Verlag GmbH.)

1581 1655 1595 1631

1604

a b c d

1700

1650

1600

1550

–1

Wavenumbers/cm

Figure 3.9 Infrared difference spectra after adsorption of lutidine on AlF3-H activated at 373 K. (a) equilibrium pressure at RT, (b) same as (a) but after evacuation at RT, (c) evacuated at 323 K, (d) evacuated at 373 K. (Reprinted with permission from [20] Copyright (2008) WileyVCH Verlag GmbH.)

Since it seems obvious that these new catalysts might be applied for other related reactions, we also tested aluminium and magnesium hydroxyfluorides for the cyclization of citronellal yielding (–)–isopolegol [21]. Some of these catalysts yielded extremely high chemo- and diasterioselectivity for (–)–isopolegol, which has not been achieved so far with other catalysts –not even with homogenous ones.

84

3.5

Functionalized Inorganic Fluorides

Oxidation Catalysis

Typically, oxidation catalysts are oxides of transition metals, the oxidation state of which changes easily. Vanadium oxide species (VOx) are well-known examples for oxidation catalysts. Their performance depends on the dispersion and coordination of VOx and on the surface properties, especially the acidity of the support on which VOx is deposited [22–26]. Thus, VOx supported by materials such as ZrO2 and TiO2, which have strong Brønsted acid sites on their surface, is very active in propane oxidation, however, its selectivity to propene, the product that is aimed for, is very low at higher conversion degrees and CO/CO2 formation dominates besides H2O because of formation of very reactive carbocations as intermediate. On the other hand, without acid sites, as in case of VOx/SiO2, very low conversion of propane occurs, obviously because the propane becomes inactive [27]. Although it is still a matter of debate in literature it is obvious that acid surface sites, beside oxygen mobility in the catalyst, are a major source of COx-formation. Surface Brønsted acid sites, in particular, are considered to activate the propane molecule in such a way that intermediate carbocations are formed which are too reactive and, hence, proceed to form COx. With the access of highly Lewis acidic HS-AlF3 it seemed plausible to use this as support for the oxygenating component, VOx because Lewis acidic supports and supports with balanced Lewis/Brønsted acidy should be able to activate the propane without formation of carbocations with their excessive reactivity, giving catalysts with improved performance. The expectation was to prevent the formation of reactive carbocationic species, which should significantly improve the selectivity toward propene and should suppress COx-formation. The sol-gel synthesis of metal fluorides provides a convenient access to VOx catalysts supported by high surface-area metal fluorides exhibiting Lewis acid and also some Brønsted acid sites (see Chapter 1). Aluminium fluoride-based VOx catalysts have been prepared from aluminium triisopropylate and vanadium(v) oxytripropoxide with both stoichiometric and understoichiometric amounts of HF/Al. In some experiments was water also added. This way, after work up and calcination, VOx catalysts have been obtained that are supported by high surface aluminium fluoride or aluminium oxyfluoride. The performance of a series of VOx/Al oxyfluoride catalysts (prepared with Al:HF ¼ 1:2) for oxidative dehydrogenation (ODH) of propane is presented in Figure 3.10. Although the yield of propene does not exceed about 12% at best, the performance of the catalysts is remarkable insofar as, for the first time, the yield of propene is higher than that of COx, with the 10% vanadium catalyst. The good performance of the aluminium fluoridesupported VOx catalysts is due to the nature of the support and not a result of the sol-gel preparation route as follows from a direct comparison of VOx/aluminium fluoride (‘VAlF’) and VOx/aluminium oxide (‘VAlO’) catalysts prepared similarly via the sol-gel route. In Figure 3.11 compares the performance of VAlF and VAlO catalysts with varied V content for propane-ODH. From Figure 3.11 (a) follows that with aluminium oxide-based catalysts high conversion rates have been obtained, even with only 10 mol% V content, however, predominantly as undesired total combustion with a low yield of propene. A rather different picture was obtained with the aluminium fluoride-based catalysts (Figure 3.11 (b)). Catalysts with low V content still gave acceptable conversion rates but the conversion resulted predominantly in the aimed propene. Only with higher V content did the conversion rate increase to nearly the same values as in the case of VAlO catalysts, however, the yield

High Surface Area Metal Fluorides as Catalysts

85

Conversion X & yield Y (%)

40 X (C3H8) Y (C3H6) Y (CO) Y (CO2)

30

20

10

0 0

5

10 15 20 V-content/mol-%

25

30

Figure 3.10 Oxidative conversion of C3H8 (X) and yields (Y) of C3H6, CO and CO2 over VOx containing aluminium oxyfluorides with different V content. (Reprinted with permission from [28] Copyright (2005) Royal Society of Chemistry.)

of propene did not increase accordingly. The superior performance of the VAlF-type catalysts led to their thorough investigation, mostly in comparison with VAlO catalysts [29]. Thus, VAlF (from Al:HF ¼ 1:3) and VAlO have, in the catalytically interesting range of up to 20 mol% V, rather similar specific surface area whereas pore volume and especially the number of acid sites (from NH3-TPD) are significantly higher for VAlF. Infrared spectroscopic investigations of adsorbed pyridine showed that the VAlF acid sites are predominantly Lewis sites. X-ray diffraction measurements showed that VAlF is amorphous and the diffraction patterns of VAlO contain only very broad reflexes of Al-O-species. The absence of VOx related diffraction patterns prove that they are incorporated in the respective Al matrices. The 27Al MAS NMR and 19F MAS NMR spectra of VAlF, depicted in Figure 3.12, resemble very closely those of HS-AlF3. The chemical shift of dF ¼ –167 ppm and the large spinning side bands are indicative of an Al in a highly distorted AlF6 environment [29]. The 51V MAS NMR spectra (Figure 3.13) and ESR spectra (Figure 3.14) of the VOxdoped aluminium fluoride and oxyfluoride systems revealed that vanadium exists there predominantly in oxidation state þIV and to a small part þV, although no signs of V2O5like species have been detected. The VAlF catalysts differ from the VAlO catalysts in their oxygen mobility, which is significantly lower for VAlF, as found with the help of 18O exchange experiments [29]. The high oxygen mobility in VAlO materials is responsible for both a high catalytic activity and predominant total oxidation to CO and CO2, i.e. low selectivity to propene. In the VAlF material there is no mobile lattice oxygen for the total oxidation available; the oxidation obviously occurs via another mechanism at the acidic surface of the catalyst, resulting predominantly in propene formation [29]. As a consequence of these investigations it seems evident that Lewis acidic metal fluorides might be interesting supports for selective oxidation catalysis because of two important advantages that they provide in

86

Functionalized Inorganic Fluorides (a)

XPropane YPropene YCOx

Conversion (X) & yield (Y) (%)

60 50 40 30 20 10 0 0

5

10

15

20

25

30

35

40

V-content (mol-%) (b) 40

XPropan

Conversion (X) & Yield (Y) (%)

35

YPropen YCOx

30 25 20 15 10 5 0 4

6

8

10

12 14 16 18 V-Content (mol-%)

20

22

24

Figure 3.11 Propane-ODH results with (a) VOx /aluminium oxide (VAlO) and (b) VOx / aluminium fluoride catalysts (VAlF) (from Al:HF ¼ 1:3) with different V contents (773 K; mcat ¼ 100 mg: N2:O2:C3H8 ¼ 20:50:30) (X) conversion; (Y) yield. (Reprinted with permission from [29] Copyright (2008) Elsevier Ltd.)

comparison to all of the metal oxide systems. These are, i) suppression or at the best prevention of any Brønsted acid sites and thus suppression of highly reactive carbocationic intermediates of the organic reactant molecule and ii) prevention of any bulk oxygen mobility in the solid catalyst, which alters the reaction path to a totally different mechanism. The usual Mars and van Krevelen mechanism does not hold for such metal fluoridebased systems. Since similar systems created with the weak Lewis acidic MgF2 system

High Surface Area Metal Fluorides as Catalysts A

87

δiso = –167 ppm

B –17 ppm

VAlF 15

VAlF 15

* 400 kHz

–400 kHz 100

0

0

*

*

–100 –200 δ (ppm)

* –300

–400

Figure 3.12 27Al MAS NMR spectrum of VAlF (15 mol% V), rot ¼ 25 kHz, central transition enlarged as insert (left), and 19F MAS NMR of VAlF, rot ¼ 30 kHz; * spinning side bands (right). (Reprinted with permission from [29] Copyright (2008) Elsevier Ltd.)

V2O5

VOF3

–615ppm

–744, –771, –783ppm

a

c 0

400kHz

–400kHz 1 MHz

0

VAlF 20 350

VAlO 20 350 –582ppm

–797ppm –519ppm

d

b 1 MHz

–1 MHz

0

–1 MHz

1 MHz

0

–1 MHz

Figure 3.13 51V MAS NMR spectra of (a) V2O5 ( rot ¼ 14 kHz), d ¼ 615 ppm; (b) VAlO (20 mol% V; calcined at 623 K) ( rot ¼ 13 kHz), d ¼ 582 ppm; (c) VOF3 ( rot ¼ 15 kHz), d1 ¼ 744 ppm, d2 ¼ 771 ppm, d3 ¼ 7835 ppm; (d) VAlF (20 mol% V; calcined at 623 K) ( rot ¼ 13 kHz), d1 ¼ 519 ppm, d2 ¼ 797 ppm. The sharp signal in (b) and (d) is of 27Al because of the very wide sweep width; an overlap of its side bands with those of V can be seen. (Reprinted with permission from [30] Copyright (2007) Elsevier Ltd.)

88

Functionalized Inorganic Fluorides

VAlF 05 VAlF 25 200

250

300

350

400

450

B0 /mT

Figure 3.14 X-band ESR spectra of VOx doped on aluminium fluoride (VAlF 15) and on aluminium oxyfluoride (VOx /AlFyOx) (Reprinted with permission from [29] Copyright (2008) Elsevier Ltd.)

gave just low conversion in these reactions, it proves that Lewis acid sites of a ‘considerable’ strength are needed for these selective oxidation reactions. Unfortunately, quantification cannot be given based on the experimental materials available. Separate oxidation experiments under liquid-phase conditions confirmed the importance of strong Lewis-acid sites on a solid metal fluoride for its performance as oxidation catalyst. HS-AlF3 and HS-Al0.9M0.1F3þd with M ¼ FeIII, MnIII, VIII, and NbIII were prepared in similar way, via sol-gel fluorination, drying and postfluorination with CHClF2 at temperatures up to 503 K, and tested as catalyst for the liquid-phase oxidation of ethylbenzene with tertbutylhydroperoxide (TBHP) [31]. Although HS-AlF3 had the least capacity for O2 chemisorption among the catalysts, it gave a good conversion rate of 42% within 6 h at 333 K; it was only inferior to the Fe and V doped catalysts but was superior to the Mn and Nb doped catalysts. The performance correlates well with the concentration (and strength) of Lewis acid sites, which were determined by the NH3-TPD and IR spectra of adsorbed pyridine. These experiments give evidence that very strong Lewis acid sites on a solid material are able to activate components of oxidation reactions.

3.6

Metal Fluoride Supported Noble Metal Catalysts

There are numerous publications in the literature about noble metal catalysed reactions – too many for us even to try to refer to them here. Noble metals have also been frequently

High Surface Area Metal Fluorides as Catalysts

89

employed for hydrodechlorination reactions of either pure chlorocarbons or chlorofluorocarbons. The synthesis of metal fluoride-supported noble metal catalysts was investigated because of the following intentions: (i) nanoscopic novel metals have attracted considerable attention from chemists, especially the catalytic community, but also among physicists. Many systems have been reported consisting of nanoscopic novel metals supported on metal oxides, very often with superior properties when compared to the classically prepared novel metal catalysts. Hence, it was expected that the fluorolytic sol-gel synthesis of metal fluorides can be readily modified to introduce finely dispersed noble metals like Pt or Pd into the high surface-area metal fluorides, instead loading the noble metal just afterwards onto the metal fluoride as has been almost done so far. It was anticipated that the noble metals might be significantly better dispersed (smaller) when their precursors are already present during the fluorolysis of the metal alkoxides so becoming in situ incorporated into the metal fluoride gel network. (ii) For several halocarbon reactions – e.g. hydrodehalogenation reactions – noble metals are the best catalysts. Such reactions involve both activation of the H-H-bond (at the noble metal site) and activation of the respective C-X-bond (presumably at the Lewis acid site). Thus, metal fluorides should be perfect supports. By using metal fluorides as support, the high thermal and chemical resistance of metal fluorides against HCl and HF can be advantageously employed in reactions where such corrosive conditions are involved (iii) Last but not least, it was intended to apply the obtained nano-nobel metal/nano-metal fluoride catalysts for some other reactions that could benefit from the high Lewis acidity of the fluoride. The noble metals have to be added as, e. g., acetylacetonates in the course of the sol-gel synthesis and after the postfluorination step necessary to bring about the full Lewis acidity of the metal fluoride, a final reduction step follows to convert the noble metal to oxidation state zero. Pd/HS-AlF3, Pd/HS-MgF2 and several other Pd/metal fluoride catalysts as well as Pt/HS-AlF3 and Pt/HS-MgF2, mostly with noble metal content around 5% (w/w), have been prepared, characterized and used as catalysts for hydrodehalogenation reactions and tested for Suzuki coupling reactions. As expected, the synthesis results in a high degree of dispersion for the noble metal (see below) whereas the properties of the respective metal fluoride are only marginally affected. From Figure 3.15 it follows that the X-ray patterns of the noble metal-containing catalysts are almost identical with those of similarly prepared metal fluorides. Thus, Pd/HS-AlF3 is as amorphous as HS-AlF3. It is interesting to note that in course of the postfluorination step the PdII signals diminish, obviously due to the reducing effect of the organic moieties liberated at higher temperatures. Such a reduction is also evident from the respective Pd 3d5/2 XPS binding energies [33]. The surface area and acidity of the noble metal catalysts are comparable although somewhat reduced compared to the respective data of the similarly prepared pure metal fluorides. The performance of these new catalysts and their advantages and disadvantages will be described below based on three different types of reactions, these being i) hydrodechlorination

90

Functionalized Inorganic Fluorides ∗-MgF2

Ο - Pd(acac)2

0

Θ -Pd

II

I

∗

∗ ∗

∗

∗ ∗

Θ

∗

d

d Ο Ο

Ο Ο Ο

c

Ο Ο

c

Ο

ΟΟ

a 40

Δ-CaF2

III

50

60

70

a 10

20

30

Δ

IV

Δ

30

40

50

60

70

-KMgF3

Δ

20

Δ

10

Δ

Δ

Δ

Intensity/a. u.

b b

Δ Δ

Ο

d

Δ

Θ

d Ο Ο Ο

ΟΟ

c

ΟΟ

c

Ο

Ο

b b a

10

20

30

40

50

60

70

a 10

2θ/degree

20

30

40

50

60

70

Figure 3.15 X-ray diffraction patterns of (I) AlF3, (II) MgF2, (III) CaF2 and (IV) KMgF3 based catalysts. (a) Pure MFx, (b) PdII/MFx-y(OR)y-precursor, (c) PdII/MFx after post-fluorination and (d) Pd0/MFx after reduction. (Reproduced from [32] by permission of Elsevier.)

of monochlorodifluoromethane, CHClF2, ii) hydrodechlorination of dichloroacetic acid, and iii) Suzuki coupling.

3.6.1

Hydrodechlorination of Monochlorodifluoromethane

After CFCs were phased out, hydrodechlorination of CFCs, especially CFC-12 (CCl2F2) and CFC-114 (CF3-CCl2F), attracted considerable attention from researchers because this was expected to yield as products HFC-32 (CH2F2) and HFC-134 (CF3CH2F), respectively, which are the most important CFC-alternatives. Early work on the hydrodechlorination of CCl2F2 reported that Pt and Pd, supported on alumina and charcoal, respectively, were very reactive catalysts [34–37]. For the hydrodechlorination of CF3-CCl2F, Pd was also found to be a more selective catalyst than Pt, although both systems gave very high yields (cf. [38–40] and references therein). It was found that especially catalyst with -AlF3 as support exhibited best catalytic properties (high conversion and high selectivity), being even better than the best reported Pd catalyst supported on char coal. Hence, we tested the Pd and Pt catalysts described above, supported on highly Lewis acidic nano metal fluorides for the hydrodechlorination of CHClF2 as given in Equation (3.17) aiming the reaction path into CH2F2.

High Surface Area Metal Fluorides as Catalysts CHClF2

+H2 –HCl

CH2F2

+H2

CH3F

–HF

+H2

CH4

91

(3.17)

–HF

Several Pd catalysts supported by AlF3, MgF2, CaF2 or KMgF3 have been prepared, characterized and used for this reaction. The supports were chosen according to their respective Lewis acidity – HS-AlF3 is the second strongest Lewis acid of all, HS-MgF2 is a weak Lewis acid, and CaF2 and, more specifically, KMgF3 are not Lewis acidic. The same order of acidity holds also for the Pd loaded systems. This can be seen from their photoacoustic IR spectra of chemisorbed pyridine of the neat HS-MFxs as well as the respective Pd-loaded ones (Figure 3.16). The surface area of the supported Pd catalysts is about 50% to 75% of the area of the corresponding neat high surface metal fluorides. This might be due to the necessary additional synthesis step, i.e. the reduction of PdII to Pd0 in H2 at 473 K.

A

BPy (1545) LPy + BPy (1490) LPy (1454)

a

Intensity/a.u.

b c d

B a b c d

1575

1550

1525 1500 1475 1450 Wavenumber/cm–1

1425

1400

Figure 3.16 FTIR photoacoustic spectra of pyridine chemisorbed at (A) MFx (a ¼ AlF3, b ¼ MgF2, c ¼ CaF2, d ¼ KMgF3) and (B) Pd0/MFx (a ¼ Pd0/AlF3, b ¼ Pd0/MgF2, c ¼ Pd0/CaF2, d ¼ Pd0/KMgF3). LPy ¼ Lewis acid sites, BPy ¼ Brønsted acid sites. (Reprinted with permission from [32] Copyright (2008) Elsevier Ltd.)

Another important characteristic of HS-MFx-supported Pd catalysts is the high Pd distribution in form of very fine particles. It has been shown by transmission electron microscopy (TEM) of freshly prepared catalyst and of a Pd0/CaF2 catalyst used for CHClF2 hydrodehalogenation reaction that the palladium is integrated as Pd-particles of 3 to 8 nm diameter into CaF2 particles of about 40 to 50 nm diameter (Figure 3.17) as originally anticipated as result of employing the fluorolytic sol-gel-synthesis route.

92

Functionalized Inorganic Fluorides

(a)

(b) 0.31 nm

0.38 nm

2 nm

50 nm (c)

(d) 0.39 nm

50 nm

2 nm

Figure 3.17 TEM images of fresh (a, b) and used (c, d) Pd 0/CaF2 catalysts for hydrodehalogenation of CHClF2. (Reprinted with permission from [32] (2008) Elsevier Ltd.)

The two marked lattice distances correspond to two different phases. Within CaF2 the {111} lattice plane distance of d111 ¼ 0.31(4) was determined and for the metallic Pd the distance of the {100} lattice plane of d100 ¼ 0.38(2). Only these two phases could be detected in these and in other fresh samples investigated, giving evidence for the absence of PdF2 or PdCl2, all Pd is in the metallic state as had already derived from the XRD patterns (see above). In case of the used catalyst (Figures 3.17 c and d) the lattice plane distance of 0.39 nm does not rule out the possible formation of PdF2 or PdO as result of the highly corrosive reaction conditions [33]. Hydrodehalogenation, the selective replacement of halogen by hydrogen on saturated halocarbons, depends on the strength of the C-X bond, the catalyst and the reaction conditions. Thus, upon hydrodehalogenation of CHClF2 C-C coupling products have been observed using Ni catalysts [41], whereas only CH3F and CH4 were obtained with silica-supported Pd [42].

High Surface Area Metal Fluorides as Catalysts

93

Using HS-AlF3 as Pd support, the extremely high Lewis acidity of the support should activate the C-X bond, thus improving catalytic activity and, possibly, selectivity. Appropriate experiments resulted, however, in a product composition not in conformance with that of Equation (3.17) [32]. There was no CH2F2 formed but unexpectedly high amounts of CHF3 (Table 3.3, entry 1). The formation of the latter gives evidence for prevailing dismutation reactions according to Equations (3.18). 2 CHClF2 ! CHCl2 F þ CHF3

(3:18a)

2 CHCl2 F ! CHCl3 þCHClF2

(3:18b)

3 CHClF3 ! 2 CHF3 þCHCl3

(3:18c)

Dismutation has to be catalysed by a Lewis acid of at least moderate strength. The very acidic HS-AlF3 support and even the moderately acidic HS-MgF2 support are not suited to boost the hydrodehalogenation. The high dismutation over these metal fluorides compared to that of HS-CaF2 can be seen from Figure 3.18. AlF3 MgF2 CaF2

Dismutation activity (%)

100

80

60

40

20

0 373

423

473

523

573

623

673

Temperature/K

Figure 3.18 Conversion of CHClF2 over neat HS-MFx. Reaction conditions: amount of catalyst ¼ 0.4 ml, CT ¼ 1 s, GHSV ¼ 3600 h1. (Reprinted with permission from [32] (2008) Elsevier Ltd.)

Consequently, with less Lewis acidic supports, such as CaF2 and KMgF3, hydrodehalogenation prevails over dismutation, as can be seen from the results in Table 3.4. Although with Pd0/CaF2 and Pd0/KMgF3 CHClF2 conversion is very low, the CH2F2 selectivity is significantly higher (around 70%) for both as compared to the 1.7% with the ‘state-of-the-art’ catalyst Pd0/C and its selectivity to CH2F2 of just ca. 5%. The slightly higher conversion (13.6 %) and especially the large amount of other byproducts are probably due to the presence of some metal salts [32].

94

Functionalized Inorganic Fluorides

In conclusion, these investigations clearly prove the potential of the fluorolytic sol-gel synthesis to result in nanoscopic, finely dispersed noble metal catalysts under in situ conditions that are the basis for catalytically highly active metal species. However, the high dismutation reactivity of these catalysts causes undesired side reactions at room temperature, which cannot at all be suppressed. On the other hand, the hydrogenation reactivity significantly reduces when neutral metal fluorides HS-CaF2 are used although particle diameter of palladium and/or platinum lies in the low nanometre scale. This clearly indicates that activation of the C-X-bond in heterogeneously catalysed hydrodehalogenation reactions is obviously the decisive reaction step.

Table 3.3 Hydrodehalogenation of CHClF2 with hydrogen over various metal fluorides supported palladium catalystsa Sr.No

Catalyst

Temp (K)

CHClF2 Conv. (%)

Product selectivity (%) CH2F2

1. 2. 3. 4. 5. 6. 7. a b

Pd0/AlF3 Pd0/MgF2 Pd0/CaF2 Pd0/KMgF3 Pt0/AlF3 Pt0/MgF2 Pd0/C

523 623 623 623 563 563 593

85.9 6.0 2.3 2.9 95.9 0.6 13.6

0 5.2 72.7 68.6 0 0 4.6

CH4

CHF3

CH3F

Otherb

32.1 26.8 27.3 31.4 32.9 48.7 55.5

31.5 39.6 0 0 44.7 51.3 33.1

28.2 6.3 0 0 2.2 0 1.8

8.2 22.2 0 0 20.1 0 5.0

Reaction conditions: amount of catalyst ¼ 0.25mL, H2/CHClF2 mole ratio ¼ 7, GHSV ¼ 3600h-1, CT ¼ 1sec. CH3Cl, CHF2-CHF2, C2H4, etc. (Reprinted with permission from [32] (2008) Elsevier Ltd.)

For hydrodehalogenation reactions in halocarbons with just one kind of halogen in it, e.g. pure fluorocarbons or chlorocarbons, respectively, these new highly Lewis acidic noble metal/MFx-catalysts should therefore be excellent candidates because dismutation as an undesired side reaction can be disregarded. Thus, starting with pure CHCl3 complete hydrodechlorination was observed over Pd0/AlF3 at 423 K whereas the Pd0/Ccatalyst gave only around 70% conversion under the same conditions, with remarkable amounts of incompletely converted mono and dichloromethanes.

3.6.2

Hydrodechlorination of Dichloroacetic Acid (DCA)

Pd/HS-AlF3 and Pd/HS-MgF2 as well the respective Pt-containing catalysts have also been employed for hydrodechlorination of dichloroacetic acid (DCA) to monochloroacetic acid (MCA) according to Equation (3.16), whereby some acetic acid (AA) is also formed. HCl2 CCOOH þ H2 ! H2 ClCCOOH þ HCl

(3:16)

This reaction is important because, upon chloroacetic acid synthesis via chlorination of acetic acid, some dichloroacetic acid is formed as byproduct. Results shown in

High Surface Area Metal Fluorides as Catalysts

95

Table 3.4 give evidence for a good performance of the Pd catalysts with about 50% yield of the aimed MCA; however, about 10% of the converted DCA is reduced to acetic acid (AA).

Table 3.4 Results of catalytic liquid phase hydrodechlorination of dichloroacetic acid* (unpublished results Diploma thesis Pratap Patil, Humboldt-University 2005) Catalyst 5% Pd/HS-AlF3 5% Pt/HS-AlF3 5% Pd/MgF2 5% Pt/MgF2

Reaction Time (h)

DCA Xa (%)

MCA Sb (%)

AA S (%)

MCA Yc (%)

3.5 8 6 8

55 11 55.5 20.6

90.7 99.3 91.6 98.7

9.3 0.7 8.4 1.3

49.9 10.9 50.8 20.3

* Reaction conditions: DCA ¼ 15 mmol (1.934 g), cat ¼ 5 wt% of DCA, Temp ¼ 170 C, H2 flow rate ¼ 20 ml/min. a conversion (X). b selectivity (S). c yield (Y).

With Pt-based catalysts the selectivity is much better than with Pd-based catalysts, however, the conversion is so low that the MCA yield is much lower than with the Pd catalysts.

3.6.3

Suzuki Coupling

Catalysts employing nano-sized Pd as catalytic active agent are widely used for Suzuki coupling reactions [43]. To make use of the advantages of heterogeneous catalysts the Pd particles have been supported on many different supports [44–48] whereby both small Pd particle size and the properties of the support were found to affect the catalytic activity. To test the suitability of metal fluoride-supported Pd0 catalysts for Suzuki coupling reactions, high surface area AlF3, MgF2 and CaF2 were tested because of their graded Lewis acidity, and additionally different syntheses for Pd0/CaF2 have been tested for cross coupling of 4-bromoanisole (4-BA) with phenylboric acid (PBA) to give 4-methoxybiphenyl (4-MBP) in liquid phase reaction [33]. The catalysts employed for and the results of the Suzuki-coupling experiments are summarized in Table 3.5. The catalytic results in Table 3.5 show the importance of the chemical nature of the support with CaF2 (3a, 3b) being superior to AlF3 (1) and MgF2 (2) even though the latter has the by far highest surface area of the metal fluoride catalysts. Further on, the advantage of the nonaqueous sol-gel fluorination synthesis over aqueous ones becomes apparent by comparing results obtained with 3a, 3b with that obtained with 4. However, all Pd0/CaF2 sol-gel syntheses (3a, 3b, 4, 6) are superior to the ‘classical’ impregnation synthesis (5) as follows from the data in Table 3.5 and in Figure 3.19, most probably because of the very small and homogeneously dispersed Pd particles obtained via sol-gel methods [33].

96

Functionalized Inorganic Fluorides

Table 3.5 Suzuki coupling of 4-bromoanisole (4-BA, I) with phenylboric acid (PBA, II) and K3PO4 (III) over Pd 0/MFx catalysts to 4-methoxybiphenyl (4-MBP) a. (Based on data taken from [33] with permission of the Royal Society of Chemistry.) Catalyst Code

Description

Preparation

1 2 3a 3b 4 5 6 7

Pd0/AlF3-HF Pd0/MgF2-HF Pd0/CaF2-HF Pd0/CaF2-R22 Pd0/CaF2-aq Pd0/CaF2-imp Pd0/Ca(OH)0.5F1.5 Pd0/C

{M(OR)x þ Pd(acac)2 þ aHF; þ HF; þ H2}c {M(OR)x þ Pd(acac)2 þ aHF; þ HF; þ H2} {M(OR)x þ Pd(acac)2 þ aHF; þ HF; þ H2} {M(OR)x þ Pd(acac)2 þ aHF; þ R22; þ H2}d {M(OR)x þ Pd(acac)2 þ HFaq; þ HF; þ H2}e {M(OR)x þ aHF; þ HF; þ Pd(acac)2; þ H2}f {M(OR)x þ Pd(acac)2 þ 1.5aHF; þ H2}g {from Aldrich, Degussa type E101 NO/W)h

SBETb/m2g-1

4-MBP-yield/%

115 225 118 59 37 62 90 –

9 11 88.1 81.9 48.3 8.1 30.6 85.9

a

Reaction conditions: I: II: III ¼ 1: 1.5: 3 (molar ratio), T ¼ 381 K, 35 mg catalyst, solvent: 12 ml of a 5:1 mixture of DMA and H2O. b Specific surface area by BETN2. c Fluorolysis of metal alkoxide and PdIIacetylacetonate in alcohol with stoichiometric amounts of aHF; post-fluorination of the dried product with gaseous HF; reduction of Pd(acac)2 with H2. Pd content of all samples adjusted to 5 wt%. d like c except: postfluorination with gaseous CHClF2. e like c except: fluorolysis with 40 % aqueous HF. f Preparation of high surface area CaF2 (S ¼ 99 m2 g-1) followed by impregnation with Pd(acac)2 in CHCl3 and H2 reduction of the dried sample. g Fluorolysis of methanolic solution of Ca(OMe)2 and Pd(acac)2 with 1.5 equivalents of aHF, followed by drying and calcinations in air at 350 C, followed by H2 reduction at 200 C. h Commercial catalyst for comparison.

Yield of 4-methoybenzophenone (%)

100

80

60

Pd/CaF2-HF(3a) Pd/CaF2-R22(3b) Pd/CaF2-Aq(4) Pd/CaF2-imp(5)

40

20

0 0

2

4 6 Reaction time/h

8

10

Figure 3.19 Catalytic activity of Pd 0/CaF2 samples for Suzuki coupling of 4-bromoanisole with phenylboric acid. (Reprinted with permission from [33] Copyright (2008) Royal Society of Chemistry.)

High Surface Area Metal Fluorides as Catalysts

97

Thus, even for the Suzuki coupling reaction, noble metals supported on nanoscopic metal fluorides obtained via the fluorolytic sol-gel-route are very active and selective although there is just a small advantage over the ‘state-of-the-art’ Pd0/Ccatalyst. It might be expected, however, that the application of this new class of noble metal catalyst would be most advantageous for reactions where at least one reaction step requires activation at a Lewis-acid site. Hence, many such reactions have still to be evaluated. For some of them these catalysts are expected to be superior to ‘classical’ oxide-based catalysts. In conclusion, in several catalysed reactions the fluorolytic sol-gel synthesis-derived nanoscopic metal fluorides have already shown strong potential. Although it is probably the case that, for many catalytic reactions, metal oxides might remain the preferred options for many reasons, there are several very promising fields for which nanoscopic metal fluoride-based catalysts are advantageous. For many halogen exchange reactions, in particular, these new catalysts exhibit unique properties.

References [1] K.O. Christe, D. A. Dixon, D. McLemore, W. W. Wilson, J. A. Sheehy and D. A. Boatz, On a quantitative scale for Lewis acidity and recent progress in polynitrogen chemistry, J. Fluorine Chem., 101, 151–153 (2000). K. O. Christe, Quantitative Lewis Acidity Scale, a Progress Report, 19th Winter Fluorine Conference, St. Petersburg (USA) 2009, Abstract 59. [2] J. Krishna Murthy, U. Groß, St. Ru¨diger, E. Kemnitz and J.M. Winfield, Sol-gel-fluorination synthesis of amorphous magnesium fluoride, J. Solid State Chem., 179, 739–746 (2006). [3] St. Ru¨diger and E. Kemnitz, The fluorolytic sol-gel route to metal fluorides – a versatile process opening a variety of application fields, Dalton Trans., 1105–1252 (2008). [4] St. Ru¨diger, G. Eltanany, U. Groß and E. Kemnitz, Real sol-gel synthesis of catalytically active aluminium fluoride, J. Sol-Gel Sci. Techn., 41, 299–311 (2007). [5] J. Krishna Murthy, U. Groß, St. Ru¨diger, V. Venkat Rao, V. Vijaya Kumar, A. Wander, C. L. Bailey, N. M. Harrison and E. Kemnitz, Aluminium chloride as a solid is not a strong Lewis acid, J. Phys. Chem. B, 110, 8314–8319 (2006). [6] H. Bozorg Zadeh, E. Kemnitz, M. Nickkho-Amiry, T. Skapin, J. M. Winfield, J. Fluorine Chem., 107 (2001) 45/52. [7] St. Ru¨diger, U. Groß, M. Feist, H. A. Prescott, S. Chandra Shekar, S. I. Troyanov and E. Kemnitz, Non-aqueous synthesis of high surface area aluminium fluoride – a mechanistic investigation, J. Mater. Chem., 15, 588–597 (2005). [8] H. Bozorg Zadeh, E. Kemnitz, M. Nickkho-Amiry, T. Skapin, G. D. Tate, J. M. Winfield, J. Fluorine Chem. 112 (2001) 225/ 32. [9] M. Nickho-Amiry, J.M. Winfield, Investigation of fluorinated surfaces by means of radiolabelled probe molecules, J. Fluorine Chem. 128 (2007) 344–352 [10] G. Eltanany, St. Ru¨diger and E. Kemnitz, Supported high surface AlF3: a very strong Lewis acid for catalytic applications, J. Mater. Chem., 18, 2268–2275 (2008). [11] K. Tanabe, T. Sumiyoshi, K. Shibata, T. Kiyoura and J. Kitagawa, A new hypothesis regarding the surface acidity of binary metal oxides, Bull. Chem. Soc. Japan 47 (1974) 1064–1066. [12] E. Kemnitz, Y. Zhu and B. Adamczyk, Enhanced Lewis Acidity by Aliovalent cation doping in metal fluorides, J. Fluorine Chem. 114 (2002) 163–170. ¨ nveren and E. Kemnitz, Mixed metal fluorides as [13] J. Krishna Murthy, U. Groß, St. Ru¨diger, E. U doped Lewis acidic catalysts systems: a comparative study involving novel high surface area metal fluorides, J. Fluorine Chem., 125, 937–949 (2004).

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[14] J. Krishna Murthy, U. Groß, St. Ru¨diger and E. Kemnitz, FeF3/MgF2: novel Lewis acidic catalyst systems, Appl. Catal., A, 278, 133–138 (2004). ¨ nveren, W. Unger and E. Kemnitz, Synthesis and [15] J. Krishna Murthy, U. Groß, St. Ru¨diger, E. U characterization of chromium(III)-doped magnesium fluoride catalysts, Appl. Catal., A, 282, 85–91 (2005). [16] H. A. Prescott, Z.-J. Li, E. Kemnitz, J. Deutsch and H. Lieske, New magnesium oxide fluorides with hydroxy groups as catalysts for Michael additions, J. Mater. Chem., 15, 4616–4628 (2005). [17] K. A. Parker and T. L. Mindt, Electrocyclic ring closure of the enols of vinyl quinones. A 2H-chromene synthesis, Org. Lett., 3, 3875–3878 (2001). [18] St. Wuttke, S. M. Coman, G. Scholz, H. Kirmse, A. Vimont, M. Daturi, S. L. M. Schroeder and E. Kemnitz, Novel sol-gel Synthesis of acidic MgF2-x(OH)x materials, Chem. Eur. J. 14, 11488/ 11499 (2008). [19] A. Hasegawa, K. Ishihara and H. Yamamoto, Trimethylsilyl pentafluorophenylbis(trifluoromethanesulfonyl)methide as a super Lewis acid catalyst for the condensation of trimethylhydroquinone with isophytol, Angew. Chem. Int. Ed., 42, 5731–5733 (2003). [20] S.M. Coman, S. Wuttke, A. Vimont, M. Daturi and E. Kemnitz, Catalytic performance of nanoscopic - AlF3 based catalysts in the synthesis of (all-rac)--tocopherol, Adv. Synth. Catal, 350, 2517–2524 (2008). [21] S.M. Coman, P. Patil, St. Wuttke and E. Kemnitz, Cyclisation of citronellal over heterogenous inorganic fluoride – highly chemo- and diasterioselective catalysts for (–)-isopulegol, Chem. Commun. 460–462 (2009). [22] E. A. Mamedov and V. Corte´s Corbera´n, Oxidative dehydrogenation of lower alkanes on vanadium oxide-based catalysts. The present state of the art and outlooks, Appl. Catal., A, 127, 1–40 (1995). [23] M.M. Bettahar, G. Costentin, L. Savary and J. C. Lavalley, On the partial oxidation of propane and propylene on mixed metal oxide catalysts, Appl. Catal., A, 145, 1–48 (1996). [24] Y. Habuta, N. Narishige, K. Okumura, N. Katada and M. Niwa, Catalytic activity and solid acidity of vanadium oxide thin layer loaded on TiO2, ZrO2, and SnO2, Catal. Today, 78, 131–138 (2003). [25] I. E. Wachs, J.-M. Jehng, G. Deo, B. M. Weckhuysen, V. V. Guliants and J. B. Benziger, In situ Raman spectroscopy studies of bulk and surface metal oxide phases during oxidation reactions, Catal. Today, 32, 47–55 (1996). [26] A. Pantazidis, A. Auroux, J.-M. Herrmann and C. Mirodatos, Role of acid–base, redox and structural properties of VMgO catalysts in the oxidative dehydrogenation of propane, Catal. Today, 32, 81–88 (1996). [27] K. Scheurell, E. Hoppe, K.-W. Brzezinka and E. Kemnitz, Bulk and surface properties of highly dispersed VOx/ZrO2, VOx/SiO2 and VOx/TiO2/SiO2 systems and their relevance for propane oxidation, J. Mater. Chem., 14, 2560–2568 (2004). [28] K. Scheurell and E. Kemnitz, Amorphous aluminium fluoride as new matrix for vanadiumcontaining catalysts, J. Mater. Chem., 15, 4845–4853 (2005). [29] K. Scheurell, G. Scholz, A. Pawlik and E. Kemnitz, VOx doped Al2O3 and AlF3 – a comparison of bulk, surface, and catalytic properties, Solid State Sci., 10, 873–883 (2008). [30] K. Scheurell, G. Scholz and E. Kemnitz, Structural study of VOx doped aluminium fluoride and aluminium oxide catalysts, J. Solid State Chem., 180, 749–758 (2007). [31] I. K. Murwani, K. Scheurell, E. Kemnitz, Liquid phase oxidation of ethylbenzene on pure and metal doped HS-AlF3, Cat. Com. 10, 227–231 (2008). [32] P. T. Patil, A. Dimitrov, H. Kirmse, W. Neumann and E. Kemnitz, Non-aqueous sol-gel synthesis, characterization and catalytic properties of metal fluoride supported palladium nanoparticles, Appl. Catal., B, 78, 80–91 (2007). [33] B. Coq, J. M. Cognion, F. Figueras and D. Tornigant, Conversion under hydrogen of dichlorodifluoromethane over supported palladium catalysts, J. Cat. 141, 21–33 (1993). [34] Z. Karpinski, K. Early and J. L. d’Itri, Catalytic Hydrodechlorination of 1,1-Dichlorotetrafluoroethane by Pd/Al2O3, J. Cat. 164, 378–386 (1996). [35] E. J. A. X. van de Sandt, A. Wiersma, M. Makkee, H. van Bekkum and J. A. Moulin, Palladium black as model catalyst in the hydrogenolysis of CCl2F2 (CFC-12) into CH2F2 (HFC-32), Appl. Cat. A 155, 59–73 (1997).

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[36] E. J.A X. van de Sandt, A. Wiersma, M. Makkee, H. van Bekkum and J.A. Moulin, Process for the selective hydrogenolysis of CCl2F2 (CFC-12) into CH2F2 (HFC-32), Cat. Today 27, 257–264 (1996). [37] Z. Karpinski, J.L. d’Itri, Hydrodechlorination of 1,1-dichlorotetrafluoroethane on supported palladium catalysts, a static-circular reactor study, Catal. Lett. 77, 135–140 (2001). [38] H. Berndt, H. Bozorg Zadeh, E. Kemnitz, M. Nickkho Amiry, M. Pohl, T. Skapin and J. M. Winfield, The properties of platinum or palladium supported on b-aluminium trifluride or magnesium dfluoride: catalysts for the hydrodechlorination of 1,1-dichlorotetrafluoroethane, Mat. Chem. 12, 3499–3507 (2002). [39] I. Kris-Murwani, E. Kemnitz, T. Skapin, M. Nickkho-Amiry and J. M. Winfield, Mechanistic investigation of the hydrodechlorination of 1,1,1,2-tetrafluorodichloroethane on metal fluoride-supported Pt and Pd, Ctal. Today 88, 153–168 (2004). [40] P. T. Patil, A. Dimitrov, J. Radnik and E. Kemnitz, Sol-gel synthesis of metal fluoride supported Pd catalysts for Suzuki coupling, J. Mater. Chem., 18, 1632–1635 (2008). [41] A. Morato, C. Alonso, F. Medina, P. Salagre, J. E. Sueiras, R. Terrado and A. Giralt, Conversion under hydrogen of dichlorodifluoromethane and chlorodifluoromethane over nickel catalysts, Appl. Catal., B, 23, 175–185 (1999). [42] R. Hina, I. Arafa and A. Masadeh, Hydrogenation of CHClF2 (CFC-22) over Pt-supported on silicabased polymethylsiloxane composite matrices, React. Kinet. Catal. Lett., 87, 191–198 (2005). [43] N. Miyaura and A. Suzuki, Palladium-catalyzed cross-coupling reactions of organoboron compounds, Chem. Rev., 95, 2457–2483 (1995). [44] D. D. Das and A. Sayari, Applications of pore-expanded mesoporous silica. 6. Novel synthesis of monodispersed supported palladium nanoparticles and their catalytic activity for Suzuki reaction, J. Catal., 246, 60–65 (2007). [45] F.-X. Felpin, T. Ayad and S. Mitra, Pd/C: An old catalyst for new applications – its use for the Suzuki–Miyaura reaction, Eur. J. Org. Chem., 2679–2690 (2006). [46] K. Shimizu, R. Maruyama, S. Komai, T. Kodama and Y. Kitayama, Pd-sepiolite catalyst for Suzuki coupling reaction in water: structural and catalytic investigations, J. Catal., 227, 202–209 (2004). [47] B. M. Choudary, S. Madhi, N. S. Chowdari, M. L. Kantam and B. Sreedhar, Layered double hydroxide supported nanopalladium catalyst for Heck-, Suzuki-, Sonogashira-, and Stille-type coupling reactions of chloroarenes, J. Am. Chem. Soc., 124, 14127–14136 (2002). [48] M.L. Kantam, K. B. S. Kumar, P. Srinivas and B. Sreedhar, Fluoroapatite-supported palladium catalyst for Suzuki and Heck coupling reactions of haloarenes, Adv. Synth. Catal., 349, 1141–1149 (2007).

4 Investigation of Surface Acidity using a Range of Probe Molecules Alexandre Vimont1, Marco Daturi1 and John M. Winfield2 1

Laboratoire Catalyse et Spectrochimie, ENSICAEN, Universite´ de Caen, CNRS, 6 Bd Mare´chal Juin, F-14050 Caen, France 2 Department of Chemistry, University of Glasgow, Glasgow G12 8QQ, UK

4.1

Introduction

Acidity and acid-base reactions are related and important features of the chemistry of molecular fluorides. In addition to the archetypal Brønsted acid anhydrous hydrogen fluoride (I), molecules in which an hydroxyl group is bonded to a high oxidation state centre (II), for example TeVI, also behave as strong Brønsted acids. The coordination numbers found most commonly in high oxidation state binary fluorides are four and six, hence important molecular Lewis acids are (III), (IV) and (V), where X can be B, Si or Sb for example, which in many situations behave as coordinatively unsaturated species. The Brønsted and Lewis acidities of these high oxidation state centres are enhanced by the high electronegativity of F; as a result, these species are used very widely in solution as reagents or molecular acidic catalysts.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Ó 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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F

F

X

H–F F

X

F F

F F

(I)

(II)

F

X

4.1.1

X

F F

(IV)

F

F

F

F

(III)

F

F

F

F

X

F F

F (V)

Setting the Scene: Metal Fluorides versus Metal Oxides

The situation is very different in the solid state. The most widely used class of solid acid reagents or catalysts is not fluorides but metal oxides, such as silicas, the transitional aluminas, aluminosilicates and high oxidation state transition metal oxides, which include titania and zirconia. One of the important properties of these materials is their relatively large specific surface areas; in contrast, the analogous metal fluorides, as conventionally prepared, have surface areas that are determined by purely geometric factors and are therefore small. Traditional solid-state synthesis methods lead to the production of such small surface area samples, whose powders are constituted by large crystallites. Conversely, soft chemistry methods, which were not applied to metal fluoride synthesis until relatively recently, allow nanosized particles to be obtained having unconstructed surfaces. Additionally, metal fluorides are extremely sensitive to traces of water. Probably due to a combination of specific surface area and susceptibility to hydrolysis, heterogeneous acid catalysts for use in large-scale fluorination processes are usually based on either chromia or -alumina, which become fluorinated in situ. In laboratory situations, metastable aluminium trifluoride phases having structures that are more open than the close-packed -AlF3 show some promise as acid catalysts.1 Other types of functionalized oxides, which have acidic properties and therefore potential catalytic value, include sulfated oxides2 (for a recent laboratory application in organofluorine chemistry see reference 3) and fluorinated clays.4 In recent work, much of which is reviewed elsewhere in this volume, surface areas of metal fluorides, for example as measured by the BET method, can be increased markedly if the synthetic method leads to aggregates of very small (sometimes called nanoscopic) particles. Such solids will exhibit properties that are very different from the conventionally prepared analogues, for example a molecular dynamics simulation of cubic nanoparticles of -AlF3 indicated the presence of structural motifs that are found also in the metastable

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phases such as b-AlF3.5 Most importantly in the context of this chapter, is the presence of a large number of four or five coordinate AlIII centres on the surface of the nanoparticles, consistent with the generation of significant Lewis acidity. In contrast, the thermodynamically stable -AlF3, as conventionally prepared, shows little or no Lewis acid character. Recent progress in the development of synthetic routes to metal fluorides, particularly those of aluminium and magnesium, which have high surface areas, has been described in earlier chapters of this volume. The significant catalytic properties of these high surface area (HS) materials have also been emphasized. Here we examine the surface acidities, both Lewis and Brønsted, of newly synthesized fluorides based on AlIII and MgII centres and pose the question ‘What is meant by a strong acid in an heterogeneous context?’

4.1.2

Some Examples of the Application of FTIR Spectroscopy to the Study of Surface Acidity in Metal Oxides

It is convenient to use metal oxides as reference points to compare the surface acidity of metal fluorides as oxides have been studied very widely and much of the methodology described for fluorides below was developed from studies of high surface area oxides. Some examples of studies made using FTIR spectroscopy serve to introduce this aspect of the subject. A method that is common in the study of surface Lewis acidity of aluminas is to use the Lewis base, pyridine (py), monitoring the IR spectrum of the chemically adsorbed species. Observable shifts in some vibrational modes, in particular 8a and 19b, occur on coordination at the surface.6,7 Adsorption of py at activated -alumina, similar structurally to the commonly studied acid, -alumina, was followed by desorption over a range of temperatures, using diffuse reflectance infrared Fourier Transform spectroscopy (DRIFTS) to monitor the surface. Observation of the py 8a mode band intensity at various temperatures was in accord with the presence of three types of Lewis acid surface sites. This conclusion was refined by studying temperature programmed desorption of py under different conditions.8 The desorption process associated with py bound to medium strength sites comprises two components, whose relative intensities are approximately 1:1. These, and the strong Lewis acid sites, can be quantified by gravimetry. In order to propose descriptions at the ‘molecular’ level for the four types of acid site, use is made of the changes observed in the –OH stretching region of -alumina as py is adsorbed and then desorbed. It is possible therefore to correlate the relative strengths of the Lewis acid surface sites, as determined from the thermal behaviour of the py 8a vibrational modes, with the temperature-dependent changes observed in the –OH stretching region. By using the proposals made previously in the literature for the various terminal –OH group environments on aluminas,9 schematic descriptions, Figure 4.1, can be generated, which take into account the interactions between coordinated py and near neighbour surface –OH groups, depicted in Figure 4.2. An example of this methodological approach is depicted in Figure 4.3, which presents spectra of pyridine resulting from adsorption at a -alumina sample, synthesized by

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Functionalized Inorganic Fluorides H O

H O ∗ AIOct

AIOct

∗ AITet

AIOct

AIOct i) Weak Lewis Acid Site

ii) Medium-Weak Lewis Acid Site

OH

∗ AITet

∗ AITet

AITet O

iii) Medium-Strong Lewis Acid Site

iv) Strong Lewis Acid Site

Figure 4.1 Schematic of the four types of surface Lewis site proposed for Z-alumina. (Reprinted with permission from D. T. Lundie, A. R. McInroy, R. Marshall, J. M. Winfield, P. Jones, C. C. Dudman, S. F. Parker, C. Mitchell and D. Lennon, J. Phys. Chem., B, 109, 11592–11601 Copyright (2005) American Chemical Society.)

H O AIOct

AIOct

H O

N

N

AIOct

AITet

AIOct i) Weak Lewis Acid Site

OH

N

N AITet

AITet

ii) Medium-Weak Lewis Acid Site

AITet

O iii) Medium-Strong Lewis Acid Site

iv) Strong Lewis Acid Site

Figure 4.2 Schematic of the pyridine complexes proposed for the four types of surface Lewis acid site in Z-alumina. (Reprinted with permission from D. T. Lundie, A. R. McInroy, R. Marshall, J. M. Winfield, P. Jones, C. C. Dudman, S. F. Parker, C. Mitchell and D. Lennon, J. Phys. Chem., B, 109, 11592–11601 Copyright (2005) American Chemical Society.)

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Rhodia (100 m2 g1). Spectrum a shows the behaviour of the sample as is, before any activation treatment. The large ‘bump’ between 3800 and 2500 cm1 is due to water adsorbed on the surface, while the features in the range 1700–1300 cm1 belong to carbonate impurities. After a thermal treatment under dioxygen, then secondary vacuum (104 Pa), only residual OH groups remain as shown by bands between 3750 and 3600 cm1 (spectrum b). Spectrum c shows the result of pyridine adsorption then evacuation on the activated sample. C–H stretches of the adsorbed molecules give rise to the features at 3200–2900 cm1,

absorbance

= 0.5 a.u.

(a) (b) (c) (d)

4000

3500

3000 2500 2000 wavenumber/cm–1

1500

= 0.2 a.u.

absorbance

(a) (b)

ν 19b ν 8a ν 8b

ν 19a

(c) (d) 1700

1650

1600 1550 1500 wavenumber/cm–1

1450

Figure 4.3 Transmission IR spectra of pyridine adsorbed on g-alumina: top, MID-IR absorption region; bottom, zoom on the ring vibration region. (a) sample before activation; (b) activated sample; (c) pyridine adsorption and evacuation at room temperature; (d) difference between c and b spectra

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whereas ring vibrations produce bands in the 1650–1400 cm1 range. For the sake of clarity, the difference between spectra recorded after pyridine adsorption and after activation is given in spectrum d. The bottom part of Figure 4.3 represents a zoom of the image presented above, in the most interesting spectral region, i.e. in the range 1700–1400 cm1. The main vibrational modes of the pyridine ring are apparent. From bands 8a it is possible to get information on the strength of the acid sites (the higher the wavenumber, the stronger the acidity), whereas the integrated intensity of the 19b furnishes a quantitative estimation of the acid sites concentration. Further details are reported in references 10 and 11. When H-bonded species are concerned, the corresponding bands for pyridine are less sensitive and closer to the frequencies observed in the liquid phase. Such a small frequency shift can make assignments ambiguous when both weak Lewis and Brønsted acid sites are expected on the surface. In such a case, the only way to prove the formation of H-bonded species is the concomitant observation of a broad (OH) band, downward shifted by several hundreds of wavenumbers from the frequency of the free OH groups. This observation is not always straightforward because this band can become very broad and overlap with (CH) bands. Moreover, for weak interactions, the 8a band is in a Fermi resonance with the 1 þ 6a combination mode, so introducing an additional difficulty to assess the acid strength of the acid centres. To circumvent these problems it was found that an accurate analysis of the whole spectrum of pyridine, including the (CH) range, may bring further insights to the acidic properties of metal oxides. Using d5-pyridine, in particular, allowed simpler (CD) spectra to be obtained, showing that both frequencies and intensities are sensitive to the adsorption mode and making it possible to distinguish among H-bonded, protonated and coordinated species.10 Particular attention was paid to the case of coordinated species, since the strength of coordination increases (CH/D) frequencies and decreases their intensities. Such variations were investigated by DFT calculations and explained by the polarization of electron density towards nitrogen to the detriment of hydrogen atoms, leading to (i) an increase of C–H bond strength and (ii), a decrease of the dipole moment derivatives for C–H stretches. These trends were correlated with the 8a and 19b frequency shifts (Figure 4.4). It appeared that the most intense (CH) band is more sensitive to small acidity differences than those in the (C¼C) range, allowing Lewis acid strength to be discriminated even for solids having a similar acidic behaviour, such as CaO, MgO and CeO2. In many circumstances, the use of a single probe molecule is not sufficient to obtain a satisfactory view of the solid surface. In such a case the combination of different probe molecules is recommended. For example, the study of surface properties of -Ga2O3, an oxide structurally related to -Al2O3, was undertaken using pyridine, dimethyl pyridine (lutidine, DMP) (for acidity), CD3CN and carbon dioxide (for basicity), and compared with the vibrations of OH groups, which can be considered as a probe molecule intrinsic to the solid itself.11 Among the different properties observed it is noteworthy that, similar to what is known for -Al2O3, the Brønsted acidity of (partially hydroxylated) -Ga2O3 was found to be very low, albeit not negligible. DMP adsorbed on -Ga2O3 and -Al2O3 showed, in both cases, a weak IR absorption band at 1649–1652 cm1 that identified the protonated DMPHþ species. However, no traces of the pyridinium ion were found when pyridine was adsorbed on either -Ga2O3 or -Al2O3. Moreover, the surface Lewis acid strength was found to be slightly smaller in -Ga2O3 than in

Investigation of Surface Acidity ν(CH)

107

ν(CD)

3105

2340

3100

2330

3095 2320

3085

(cm–1)

(cm–1)

3090

3080 3075

2310 2300 2290

3070 2280

3065 3060 1437 1442 1447 1452 1457 1462 1467 ν19b (cm–1)

2270 1290 1300 1310 1320 1330 1340 1350 ν19b (cm–1)

Figure 4.4 Frequency of the most intense n(CH/D) band versus the frequency of the n19b mode of h5-pyridine (left) and d5-pyridine (right) coordinated on various metal oxides. (Reprinted with permission from A. Travert, A. Vimont, A. Sahibed-Dine, M. Daturi, J.-C. Lavalley, Appl. Catal. A: General, 307, 98–107 Copyright (2006) Elsevier Ltd.)

-Al2O3, due to the smaller polarizing power of coordinatively unsaturated GaIV ions, which have a smaller charge/radius ratio than Al3þ. In contrast, the concentration of Lewis acid sites (detected by adsorbed pyridine and DMP) was found to be higher in -Ga2O3, thanks to the higher tetrahedral preference of the Ga3þ ion. Both -Ga2O3 and -Al2O3 showed a similar surface basicity toward adsorbed carbon dioxide and acetonitrile. However, the surface concentration of basic O2 ions was found to be larger for gallium oxide, as deduced from the larger amount of acetamide species formed upon adsorption of acetonitrile.

4.1.3

A Preview

Having set the scene with some descriptions of surface chemistry of some metal oxides, we shall turn to the ambiguities in surface acidity determinations compared with the now rather precise acidity determinations that are possible for molecular fluorides. A possible benchmark for solid fluoride Lewis acids is aluminium chlorofluoride. Experimental methodology is then described, both the familiar FTIR method, developed directly from oxide studies (cf. § 4.1.2) and the less familiar use of chlorine-36 labelled anhydrous hydrogen chloride, used either directly or generated at a surface by dehydrochlorination of tert-butyl chloride, also labelled with [36Cl]. The use of both techniques for the investigation of various recently synthesized AlIII and MgII fluorides is described separately; finally a multi-probe approach, involving both techniques, is used to investigate the rather complex surface acidity of the hexagonal tungsten bronze aluminium hydroxy fluoride.

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4.2

Characterization of Acidity on a Surface: Contrasts with Molecular Fluorides

One of the driving forces behind recent developments in the synthesis and characterization of new acids is the requirement to replace mineral acids by reagents and catalysts that will lead to reduction in waste, in other words to an improved atom economy. Thus, HCl used in the synthesis of diamino diphenyl methane, an intermediate in polyurethane production, can be replaced by a silico-aluminate catalyst in which the acid sites of the zeolite are accessible through the external surface.12 Less conventionally, simple aromatic hydrocarbons can be nitrated using 60–70% nitric acid in the presence of the strong Lewis acids, lanthanide-perfluoroalkane sulfonic acid salts, such as the ytterbium(III) compounds (VI) and (VII).13 Yb(OSO2CF3)3

Yb(OSO2C8F17)3

Yb[C(SO2RF)3]3

Hf[N(SO2C8F17)2]4

(VI)

(VII)

(VIII)

(IX)

The class of molecular Lewis acids, encompassing f and d block metal derivatives of the strong Brønsted acids, HC(SO2RF)3 and HN(SO2RF)2, for example (VIII) and (IX), finds extensive use in laboratory synthesis, particularly Friedel Crafts acylation14 and other acid catalysed reactions of organics.15 Separation techniques are crucial, biphasic fluorous methods often being used where the RF group is C8 or greater. Supporting the Lewis acids on a polymer is another possibility. Although these catalysts have been very successful in small-scale syntheses, the environmental persistence of molecules containing long chain perfluorocarbon groups would be a disadvantage for large scale processes. Cost is another factor that favours the use of oxides and possibly binary fluorides as acids.

4.2.1

Molecular Brønsted and Molecular Lewis Acids

Molecular Brønsted acids have been studied widely in a variety of liquid media following the pioneering work of R. J. Gillespie and his school.16 Fluorine-containing acids are used widely and provide some of the strongest Brønsted acids known. Quantitatively, acidity is measured by determination of the Hammett acidity function, Ho (X). There are several spectroscopically based methods for its determination, making use of series of aromatic hydrocarbon derivative weak bases, B, which are partially protonated in solution.17 B + H+

BH+ Ho = pKBH+ – log[BH+] / [BH] (X)

The area of super acidic media, defined as media with Ho values more negative than that of 100% H2SO4 (–11.9) has been reviewed fairly recently, with emphasis on fluorinecontaining acids.18 Unlike the situation for Brønsted acidity, manifest by complete or partial proton transfer from acid to base, Lewis acidity is a more general concept. It encompasses the BrønstedLowry definition and thus there is no universal scale (in contrast to the Hammett acidity

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109

scale). Many alternative scales have been devised. Acidity and basicity cannot be related to a ‘universal’ e pair donor and relative orders of Lewis acidity will depend on the identity of the reference base used. Ideally, since an acid-base interaction is equivalent to bond energy, strengths of Lewis acids versus a chosen base should be determined by calorimetry, or an equivalent method that yields a thermodynamic quantity. The experiment should be carried out with isolated molecules, usually translated as being in the gas phase or in a nonsolvating solvent (if this latter is not a contradiction in terms). Relatively few systems have been studied in this way, hence the widespread use of computations using DFT methods and of spectroscopic methods leading to relative shifts (IR or NMR are the most popular); these types of approaches to the relative Lewis acidities of the heavier Group 3 halides have been compared.19 Comparisons among molecular Lewis acid fluorides are often made on the basis of the F ion affinities of the isolated molecules, for example using the pF scale.20 The ordering of Lewis acidities is based largely on computational work, the values being referenced to the experimentally determined gas phase F ion affinity of OCF2 (XI). OCF2 (g) + F– (g)

OCF3–(g) (XI)

From these computations and related work,21, 22 it has been established quantitatively that antimony pentafluoride and its oligomers are benchmarks for molecular, strong Lewis acids.

4.2.2

A Possible Benchmark for Solid Metal Fluoride, Lewis Acids: Aluminium Chlorofluoride

The process of making comparisons among solid Lewis or Brønsted acids is more complicated than for the molecular acids illustrated above. For example the use of Hammett indicators to determine Brønsted acidity in sulfated oxides has proved to be controversial;2 the values obtained can be misleading because heterogeneous conditions are very different from those in a nominally ideal solution. Not only are the intrinsic strengths of the different types of surface acid important but surface site densities and morphology (the latter could be in conflict with the steric requirements of the species used as a probe) may be crucial in determining the apparent effectiveness of a solid acid. In the context of fluorine chemistry, replacement of the strong Lewis acid, liquid SbF5, by a strong, solid Lewis acid, which is easily separable and thus reusable, would result in significant waste and energy reductions. The hypothetical, monomeric molecule AlF3 has a high pF value from the DFT computation20 and for this reason, development of new aluminium fluorides as potentially strong Lewis acids has attracted particular attention. Solid aluminium chlorofluoride (ACF)19, 23 was an early candidate, notwithstanding its extremely hygroscopic nature, and in some types of reaction its behaviour resembles that of liquid SbF5. For example, both compounds behave as catalysts for the electrophilic addition of fluoroalkanes to perfluoroolefins, although the former is the catalyst of choice using halopolyfluoroalkanes,24 while the latter is more effective using hydrofluoroalkanes.25

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Functionalized Inorganic Fluorides

Aluminium chlorofluoride is an amorphous solid of high specific surface area. It is prepared by halogen exchange between aluminium trichloride and a chlorofluorocarbon under conditions where exchange is almost complete; its composition is AlClxF3x, x ¼ approximately 0.05–0.3.26 A structure has been proposed from data analyses involving several physical methods.19,27 Chlorine is an integral part of the environment about octahedrally coordinated AlIII (there is no evidence in this composition range for an AlIII chloride phase); long range order among the various structural motifs is absent; those proposed are (XII): [Al(mF)5/2F] 3[Al(m3Cl)1/3(mF)5/2] n[Al(mF)6/2] : n is variable. (XII) Although there is no specific information available regarding the exact composition of the surface, it is logical to assume that coordinatively unsaturated AlIII sites are present having a variety of F, Cl environments. The compound ACF is the benchmark for the [36Cl]-radiotracer approach to Lewis acidity which is described below. The compound has therefore been described in some detail.

4.3 4.3.1

Experimental Methodology FTIR Spectroscopy

The in situ characterization of a solid surface by FT-IR is a well known and current method of investigation in a number of laboratories. Nevertheless, each team of scientists may use different analysis techniques. In our case we choose the transmission technique, so that samples are pressed (109 Pa) into self-supported discs (2 cm2 area, 7–10 mg cm2). They are placed in a home-made quartz cell equipped with KBr windows (Figure 4.5). A movable quartz sample holder permits adjustment of the pellet in the infrared beam for spectral acquisition and its displacement into a furnace at the top of the cell for thermal treatments. The cell is connected to a vacuum line for evacuation, calcination steps (residual pressure ¼ 103–104 Pa) and for the introduction of probe molecules (in gas or vapour phase) into the infrared cell. Spectra are recorded at room temperature. In adsorption experiments where probe molecules having weak interactions are used (such as CO), the temperature of the pellet is decreased to about 100 K by cooling the sample holder with liquid N2 after quenching the sample from the thermal treatment temperature. The addition of accurately measured increments of probe molecules in the cell (a typical increment corresponds to 100 mmol of probe per g of material) is possible via a calibrated volume (1.50 cm3) connected to a pressure gauge for the control of the probe pressure (1–104 Pa range). The probe pressure inside the IR cell is monitored by another pressure gauge (1–103 Pa range). Transmission IR spectra are recorded in the 5600–500 cm1 range, at 4 cm1 resolution, on a Nicolet Nexus spectrometer equipped with an extended KBr beam splitting device and a mercury cadmium telluride (MCT) cryodetector. Prior to the adsorption of probe molecules, the samples are activated at different temperatures (ranging from 373 to 873 K, depending on the material and impurity stabilities) overnight.

Investigation of Surface Acidity

111

Gauge pressure location

Wafer holder hook (top position)

Glass squared pipe

Wafer holder

To vacuum apparatus

Sample Heater element

IR

KBr windows

Figure 4.5 Schematic representation and photograph of the in situ IR cell. Figure kindly provided by Dr. P. Bazin, LCS, Caen, France

The principle of probe molecule use to investigate a solid surface is conceptually simple; a molecule (whose vibrational modes are well known) in the gas or vapour phase is sent onto the surface to be analysed. The modification of the probe vibrations due to the interaction(s) will reflect the state of the surface itself, i.e. the physico-chemical properties of the solid, the concentration of the adsorption sites, their strength, the mode of coordination, etc. The use and the choice of appropriate probe molecules have been detailed in numerous papers and in particular in reference 28, which should be consulted for further information.

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Functionalized Inorganic Fluorides

4.3.2

4.3.2.1

Characteristic Reactions and the Detection of Adsorbed Species by a Radiotracer Method State of the Art

A ‘probe reaction’ can be used to make comparative statements about the nature of different acidic surfaces. This is an approach of long standing and there are many applications described in the literature. Two examples will suffice to introduce the topic, since other chapters in this volume deal specifically with reactions that are catalysed by solid metal fluorides and related compounds. The dehydration of isopropanol to give propene has been used in making comparisons of solid acidic oxides2 and the isomerization of 1,2-dibromohexafluoropropane, CBrF2CBrFCF3, to its 2,2-dibromo-isomer, CF3CBr2CF3, is a reaction diagnostic for very strong solid Lewis acids such as ACF.19,23 In the FTIR approach, described above, there is a reasonably direct link from spectroscopic data to relative acidities and, at least for solid oxides, a substantial literature linking spectroscopic properties of probe molecules to acid site strengths. In contrast, the use of probe reactions to make statements about acidity has a smaller literature and a linkage that is less direct. However, providing comparative studies are undertaken under identical or near-identical conditions, the probe reaction approach can provide useful comparative data on catalyst activities and is complementary to the more widely used FTIR method. 4.3.2.2

Anhydrous Hydrogen Chloride as a Probe

The focus here will be on anhydrous HCl, since, (i) it is a strong Brønsted acid (at least in water), (ii) in principle it can behave as a Brønsted acid by dissociative adsorption at a surface that contains oxide (hydroxide) functioning as a Brønsted base and (iii) it could function as a weak Lewis base as a result of associative adsorption at a surface that contains (strong) Lewis acid sites. From this analysis, it would be expected that the behaviour of HCl towards fluorides would be different from its behaviour towards oxides. This comparison is less clear cut however if oxyfluorides and partially hydrolysed (hydroxylated or hydrated) fluoride surfaces are included. Schematic representations of dissociatively adsorbed HCl at medium strong and weak Lewis acid sites of -alumina are shown in Figures 4.6 and 4.7. These are based on spectroscopic observations. The three possible modes of adsorption of anhydrous HCl at a metal fluoride surface, physically adsorbed, associatively adsorbed and dissociatively adsorbed (the latter two cases are chemical adsorption) are represented diagrammatically in Figure 4.8. Anhydrous HCl can be used as a probe either by direct exposure to a surface or it may be generated in situ by a dehydrochlorination reaction (XIII). H CI + HCI (XIII)

Investigation of Surface Acidity

113

CI OH

OH

HCI

∗ AlTet

AlTet

AlTet

AlTet

O

+ OH(ads)

O – H2O CI

Al

∗

AlTet O

Figure 4.6 Proposed adsorption of anhydrous HCl at a medium strong Lewis acid site of Z-alumina and the effect of subsequent H2O loss. (Reprinted with permission from A. R. McInroy, D. T. Lundie, J. M. Winfield, C. C. Dudman, P. Jones, S. F. Parker and D. Lennon, Catal. Today, 114, 403–411 Copyright (2006) Elsevier Ltd.)

H O

O ∗ AIOct

AIOct

HCI

AIOct

AIOct

CI

AIOct AIOct

Figure 4.7 Schematic representation of HCl dissociatively adsorbed at a coordinatively unsaturated AlIII site of Z-alumina approximating to the weak Lewis acid type. (Reprinted with permission from A. R. McInroy, D. T. Lundie, J. M. Winfield, C. C. Dudman, P. Jones, S. F. Parker and D. Lennon, Catal. Today, 114, 403–411 Copyright (2006) Elsevier Ltd.)

HCI (g) HCI

HCI

HCI

HCI

HCI

HCI

HCI

HCI

HCI

and/or

CI

CI H

phsical adsorption, weakly bound

HCI

H

associative or dissociative chemical adsorption, strongly bound

Figure 4.8 Three possible modes of adsorption for anhydrous HCl at metal fluorides. (Reprinted with permission from M. Nickkho-Amiry and J. M. Winfield, J. Fluorine Chem., 128, 344–352 Copyright (2007) Elsevier Ltd.)

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Functionalized Inorganic Fluorides

Dehydrochlorination of hydrochlorocarbons in the vapour phase occurs under homogeneous flow conditions at moderate temperatures.29, 30, 31, 32 Reaction occurs either via a radical chain mechanism, for chlorinated ethanes,29 or via a unimolecular pathway, for t-butyl chloride,30 and for two chloropropanes.31 1,1,1-Trichloroethane undergoes dehydrochlorination by both pathways.32 In the presence of a clean Pyrex surface a, faster, heterogeneous pathway becomes available, for example for ButCl,30 and, more generally, heterogeneous processes, usually observed > 400 K, are the norm for hydrochlorocarbons in the presence of Lewis acidic oxides, such as CH3CCl3 at -alumina,33 CH3CHClCH3 at several acidic oxides and dichloropropane isomers at silica-alumina.34 Heterogeneous, Lewis acid catalysed dehydrochlorination occurs also as one reaction in the complex oxychlorination process catalysed either by CuCl2 supported on -alumina35 or by halide melts.36 In striking contrast to the dehydrochlorination of hydrochlorocarbons at oxides, observed above room temperature, is the room temperature dehydrochlorination of CH3CCl3 in the presence of resublimed solid aluminium trichloride,37 -alumina that has been chlorinated using CCl4 in order to enhance surface Lewis acidity38 or -alumina previously fluorinated at room temperature using sulfur tetrafluoride.39 Dehydrochlorination is accompanied by the formation on the surface of unsaturated oligomers derived from CH2¼CCl2. These are dark red or purple, therefore even a trace reaction can be easily detected. The process is inhibited by traces of water. It should be noted however that not all Lewis acid fluorides give rise to this phenomenon. The behaviour of mixed metal fluorides, Lewis acid composites formed from -CrF3,3H2O and FeF3, resembles that of -alumina rather than the fluorinated -alumina described above.40 Rather surprisingly, ButCl undergoes dehydrochlorination at room temperature in the presence of either SF4-fluorinated -alumina or b-aluminium trifluoride.41 As a result, this species, together with HCl, have been used as probes of surface acidity. The events that are possible at a fluorinated surface following its exposure to ButCl vapour at room temperature are shown in Figure 4.9.

4.3.2.3

Direct Monitoring of Metal Fluoride Surfaces using Chlorine-36 Labelled HCl

Anhydrous HCl can be detected readily in the gas phase above the surface due to its characteristic IR spectrum. By using chlorine-36 labelling ([36Cl] is a b emitting isotope, t1/2 ¼ 3.01 105 y)42 it can be detected as a surface species, physically or chemically adsorbed as indicated in Figure 4.8. Direct radiochemical monitoring of a surface is a very sensitive and well established method, which was developed originally for the study of [14C]-labelled hydrocarbons at supported metal heterogeneous catalysts.43 It has been used extensively in Glasgow for the study of fluoride and fluorinated oxide catalysts.44 The apparatus used is shown in Figure 4.10. The evacuable Pyrex counting cell, which is connected to a gas-handling system, contains two end-window Geiger-Mu¨ller counting tubes, mounted to ensure identical counting geometry, and a moveable Pyrex boat. The use of the apparatus to demonstrate adsorption is illustrated by an experiment in which aliquots of [36Cl]-anhydrous hydrogen chloride are added successively to the cell, which contains a thinly spread layer of the

Investigation of Surface Acidity CH3

H3C

H3C H3C

115

CH3

CH3 Cl

Cl ∗

H

H3C H3C

+

∗

(C4H8)n

H+

+

Cl–

H

CxHy species

*

*

H3C

H HCl

F-alumina (SF4) H CH3 + other products

β-AlF3

Figure 4.9 Reactions and other events that are possible following the adsorption of ButCl at a metal fluoride surface. Surface adsorbed species are indicated by a broken line to an asterisk. (Reprinted with permission from M. Nickkho-Amiry and J. M. Winfield, J. Fluorine Chem., 128, 344–352 Copyright (2007) Elsevier Ltd.)

benchmark compound, aluminium chlorofluoride, ACF (see Section 4.2.2 above).44 The procedure is as follows. A sample of ACF is dropped into the cell in vacuo via a Pyrex ampoule (A in Figure 4.10), which is attached vertically to the counting cell. Prior to its use, the flamed out ampoule has been loaded with the solid in a glove box. After it is added to the boat, the sample (250–500 mg) is spread as thinly as possible on the base of the boat. The latter is positioned, using a magnet and soft iron bar encased in the boat handle, directly below one of the counters (C1 in Figure 4.10). The arrangement and procedure enable the requirement for an infinitely thin layer of solid to be located directly below one counter, to be approximately fulfilled. The solid and the cell are outgassed thoroughly and then a measured pressure of [36Cl]-labelled anhydrous HCl is expanded into the cell from the gas-handling manifold (B in Figure 4.10). The cell is isolated and its contents allowed to equilibrate for 15 min. Counts from each G.-M. counter are recorded simultaneously on two scaler timers, ideally for a sufficient time to accumulate a substantial number of counts. Since the counting error is the square root of the count number,42 to achieve a relative error of 1% requires 104 counts to be accumulated. This is not always possible and a counting time of 500 s is an acceptable compromise. After counting, the vapour is removed from the cell at room temperature by condensation at 99 K and the sequence: admission of gas, equilibration and counting, repeated. In the experiments described here,

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Functionalized Inorganic Fluorides

A

B C1

C2

solid counting apparatus

Figure 4.10 A schematic of the Geiger-Mu¨ller direct monitoring process. C1 and C2 are two intercalibrated end window Geiger-Mu¨ller counting tubes, A is an evacuable ampoule, originally containing the solid fluoride; B is an evacuable ampoule from which the labelled volatile probe compound is dispensed. (Reprinted with permission from M. Nickkho-Amiry and J. M. Winfield, J. Fluorine Chem., 128, 344–352 Copyright (2007) Elsevier Ltd.)

usually eight or nine aliquots were used for each series. After removal of the last aliquot, the outputs counts from both counters are recorded. The cell is left evacuated for at least 18 h, usually with intermittent pumping, before additional counts are recorded. Background counts are obtained before and after each experiment and counters are changed when their contamination is suspected. Counters are intercalibrated regularly using a range of pressures of the labelled vapour being used. The intercalibration ratio is the gradient of the linear relationship between counts from the two counters. Experimental data obtained from exposures of H36Cl aliquots to ACF are shown in Figure 4.11. The count from C1, Figure 4.11 (a), reflects the sum of the surface count and that from the cone of gas between the counter end window and the surface. A build up of activity on the surface is indicated, suggesting that as gas and presumably weakly adsorbed species are removed between each addition, some of the H36Cl is strongly bound. The relatively large specific surface area (approximately 100 m2 g1) of ACF is ideal for this technique and results in a relatively high degree of precision. The counts from C2 reflect the H36Cl in the gas phase; their magnitudes are far smaller than those counted by C1 and accordingly the precision is lower. Intercalibration factors are usually close to 1, therefore the count number relationship with C2 data is very similar to the intercalibrated values (C2(ic), shown in Figure 4.11 (b). The surface count is the difference between C1 and C2(ic) and is shown in Figure 4.11 (c). The relation of this variation in surface count across the range of

Investigation of Surface Acidity 4000

13 500 intercalibrated C2 500 s count

(a) 13 000 C1 500 s count

117

12 500 12 000 11 500 11 000

3900

(b)

3800 3700 3600 3500 3400 3300 3200 3100 3000

10 500 0

2

4

6 count No.

8

0

10

2

4

6 count No.

8

10

C1- C2(ic) 500 s surface count

9500 (c) 9000 8500 8000 7500 7000 0

2

4

6 count No.

8

10

Figure 4.11 Sequential exposures of H36Cl aliquots (2.0 kPa) to aluminium chlorofluoride (ACF). (a) C1 counts surface þ the volume of gas directly above; C2 counts an equivalent gas volume; (b) C2(ic) is the intercalibrated count from C2. (c) The derived surface count, obtained from C1 – C2(ic). (Reprinted with permission from M. Nickkho-Amiry and J. M. Winfield, J. Fluorine Chem., 128, 344–352 Copyright (2007) Elsevier Ltd.)

additions, together with the analogous relationship when aliquots of [36Cl]-ButCl vapour are exposed to ACF, are discussed below.

4.4

Experimental Studies of Surface Acidity

The fluoride materials to be highlighted are, firstly, the high surface area (HS) fluorides of aluminium(III)45 and magnesium(II)46 prepared by the sol-gel route and, secondly, derivatives of b-aluminium trifluoride, particularly hydroxy, fluorides prepared by a microwave-assisted solventothermal reaction in aqueous-organic HF media.47,48,49 In both cases the synthetic routes are described in detail in other chapters of this volume. The types of information obtained from FTIR spectroscopy and by using [36Cl]-labelling are described in turn. Finally the acidity of an aluminium hydroxy fluoride having the hexagonal tungsten bronze (HTB) structure,50 which was prepared by the

118

Functionalized Inorganic Fluorides

microwave-promoted solventothermal route referred to above,47 will be described using data from both techniques.

4.4.1

Using FTIR Spectroscopy

Despite the great interest in fluoride compounds for catalysis, their surface properties have been little studied up to now, compared with the corresponding metal oxides. Moreover, the published studies deal mainly with material having specific surface areas less than 20 m2 g1. Recently we have investigated well crystallized and amorphous iron, chromium and aluminum fluorides, all presenting a high specific surface area, up to 300 m2 g1. We characterized first crystalline hexagonal HTB structure fluorides, synthesized in Bordeaux. These solids are constituted by MF6 octahedra linked by corners.50 They can ˚ in be considered as microporous materials, having monodirectional channels of 3–4 A diameter, which can contain water or ammonia. Their thermal stability varies between 423 and 623 K according to the nature of the cation (Fe, Cr, Al or Ga). In addition, we have studied materials having the pyrochlore structures, synthesized in Bordeaux, as well as amorphous compounds of a new type, synthesized using a sol-gel method in Berlin and having a very high specific surface area. The identities of the compounds to be discussed below are summarized in Table 4.1. Table 4.1 Formulation, method of synthesis and specific surface area of the samples synthesized by the ICMCB team in Bordeaux and the HU-Berlin Institut fu¨r Chemie team Composition and structure

Synthesis

HTB-AlF2.2(OH)0.8a HTB AlF3-x(OH)x HS-AlF3 post fluorinated (F2) HS-AlF3 post fluorinated (CFC) Pyrochlore AlF1.8(OH)1.2

nitrate route nitrate (microwave activation) alkoxide route (microwave activation) alkoxide route (sol-gel)

a

alkoxide route (microwave activation)

Specific surface area/m2 g1

Reference

15 82 330

51, 52 47 48

180–310

53 (see also 45) 49

137

Study extended to MF3-x(OH)x , with M ¼ Fe, Cr, Ga.

4.4.1.1

Chemical Impurities and Hydroxyl Groups

Fluorides often present surface impurities as residuals from synthesis.47–53 Their identification can elucidate the reaction mechanisms; in particular, it has been demonstrated by IR spectroscopy that aluminium hydroxy fluorides having the HTB structure contain coordinated or protonated ammonia.47 Curiously, their synthesis involves a nitrate Al(NO3)3,9H2O precursor and ammonia is absent as a reagent. Therefore we have deduced that precursor decomposition implies nitrate reduction to ammonia by isopropanol during the reaction, thus explaining the presence of both ammonia inside the material and acetone

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119

in the filtrate residual. This result has allowed a better understanding of the microwave assisted synthesis.47 In general, the main impurity is residual water, which can generate OH groups. Their localization, inside the bulk or at the surface, is important to understand solid reactivity. A preliminary study of FeIII, AlIII or CrIII HTB compounds51 has shown the presence of a sharp and intense (OH) band, assigned to structural M-OH groups.52 The use of deuterated probe molecules having different sizes has shown that most of these OH groups are not sensitive to H/D exchange; the steric hindrance of the probe molecule is thus greater than the channel size. Therefore hydroxyls are localized inside the channels; their formation is likely to be due to a F/OH substitution induced by water during the synthesis step. This replacement matches well with a fluorine content significantly less than that expected for pure MF3 fluorides, so that it is more appropriate to define such solids as hydroxyfluorides having a MF3x(OH)x formula.51, 52 More recently, a structural study dealing with aluminum based compounds47 has shown that OH groups bridged between two aluminum atoms are localised in particular crystallographic positions, i.e. F1 and F2 sites (Figure 4.12).

Figure 4.12 A view along the c axis of the AlF3 HTB compound. (Reprinted with permission from D. Dambournet, A. Demourgues, C. Martineau, S. Pechev, J. Lhoste, J. Majimel, A. Vimont, J.-C. Lavalley, C. Legein, J.-Y. Buzare´, F. Fayon and A. Tressaud, Chem. Mater., 20, 1459–1469 Copyright (2008) American Chemical Society.)

Infrared spectra of the compounds indicate that the frequency of (OH) bands is relatively low whatever the identity of the cation (AlIII, FeIII, CrIII), compared with those of the corresponding metal oxides, whereas the position of the (OH) mode is unexpectedly high. Analogous variations in (OH) and (OH) wavenumber values have been observed for bridged hydroxyls in zeolites having a small pore size. Thus, the unusual

120

Functionalized Inorganic Fluorides

position for the stretching and bending modes for hydroxyls present in the tunnels of the HTB framework can be explained by the bridged conformation of the hydroxyls in a confined environment.52 For the aluminium compounds we have reported a (OH) component situated at a higher wavenumber (3680 cm1) with respect to that observed for the Al-OH inside the channels (3665 cm1).47,54 This feature is sensitive to the H/D exchange via D2O, as well as to the adsorption of large probe molecules such as lutidine (DMP) or pyridine, therefore we have assigned it to Al-OH groups on the external surface of crystallites; the absence of a confinement effect would be the reason for its higher frequency. This phenomenon is not specific to the HTB structure; the pyrochlore aluminum hydroxy-fluorides (microporous compounds having a related structure) also present a different (OH) frequency between hydroxyls inside the channels (3673 cm1) and those exposed at the crystallite external surface (3720 cm1).49 An unpublished theoretical modeling of structure and surface of partially hydroxylated HTB aluminum fluorides, carried out by the STFC Daresbury Laboratory team, has confirmed such vibrational behaviour. 4.4.1.2

Brønsted Acidity of Metal Fluorides

Brønsted acidity of a series of HTB materials has been studied by probe molecule adsorption (ammonia, pyridine, lutidine, CO) monitored using infrared spectroscopy. We have observed that the results depend on the probe used and on the activation conditions. The OH groups present in the channels are accessible to ammonia only for Fe and Fe/Cr compounds, while Al samples are unaffected by probe substitution.52 ˚ ) is smaller than Structural analysis confirms that the channel size for aluminum (2.42 A ˚ ), corresponding to the observations made using NH3 and consistent with the for iron (2.7 A greater sizes of pyridine or lutidine. Through ammonium cation formation we have shown that Brønsted acidity is correlated with the presence of such OH groups inside the micropores. Another source of acidity would be the presence of HF inside the channels. Unfortunately it was not possible to confirm such an hypothesis by the observation of IR bands corresponding to this species.52 However, the study of hydroxyl band intensity variation versus adsorbed ammonia quantity for the iron and chromium compounds suggests an additional source of Brønsted acidity, because the ammonium species number deduced is greater than the hydroxyl number. Consequently, it would be interesting to adsorb CO on Fe- or Cr-based compounds, because the HF. . . CO complex is claimed to give a characteristic band at 2172 cm1.55 A more complete study of an aluminum HTB compound, stoichiometry AlF2.6(OH)0.4 and a greater specific surface area (82 m2 g1), shows that external OH groups have sufficient acidity to protonate lutidine, but not pyridine.54 Complementary CO adsorption experiments at low temperatures have confirmed the acidic character of those hydroxyls, presenting a remarkable homogeneity in acidity strength, as confirmed by the D(OH) ¼ 3680 ! 3520 cm1 shift, correlated with D(CO) ¼ 160 cm1, corresponding to a (CO) frequency observed at 2173 cm1 (band E in Figure 4.13). This acidity is definitively lower than those for acidic solids such as zeolites, but clearly greater than that reported for OH groups in aluminas. Thus the fluorine effect on the creation of strong Brønsted sites is established.54 Another Brønsted acidity source is the presence of water on the surface of crystalline or amorphous solids. Thus water addition on the HTB compounds transforms coordinated

Investigation of Surface Acidity

121

3525 0.04

0.1 (d)

3685 (g)

(g) (d)

3700 3695 3690 3685 –1 /cm

3580 3540

3500

/cm

–1

E 2173

abs 0.01

(g)

F 2166 D 2183

(a)

Physisorbed species.

B C 2215 2200 A 2235

(f)

2240

2220

2200

2180

2160

2140

2120

2100

–1

wavenumber/cm

2240

2220

2200

2180

2160

2140

2120

2100

–1

wavenumber/cm

Figure 4.13 Left: IR spectra recorded at 100 K after introduction of increasing CO doses at 100 K: 10 mmol g1 (a), 44, 71, 88, 130, 300 mmol g1 (b to f, respectively) then an equilibrium pressure (665 Pa) (g). Right: deconvolution of the n(CO) band envelope on the spectrum (f). Inset: spectral region of the perturbed n(OH) bands. (Reprinted with permission from D. Dambournet, H. Leclerc, A. Vimont, J.-C. Lavalley, M. Nickkho-Amiry, M. Daturi and J. M. Winfield, Phys. Chem., Chem. Phys., 11, 1369–1379 Copyright (2009) PCCP Owner Societies.)

ammonia species into ammonium. Similarly, water addition transforms Lewis acid sites into Brønsted sites on activated amorphous compounds. Due to the strong Lewis acidity, water is strongly coordinated on fluorides and generates strong Brønsted sites, as evidenced by CO adsorption.53 On metal oxides, it is well known that water can transform Lewis into Brønsted acidity by dissociative water adsorption on the Mþ – O acid-base pairs, with the consequent creation of M-OH acidic groups. The absence of sufficiently strong basic sites on fluorides, as shown by an unpublished study of propyne adsorption, inhibits dissociative adsorption of water on such materials. 4.4.1.3

Lewis Acidity of Metal Fluorides

The studies performed on a wide range of compounds as indicated in Table 4.1, have allowed three general factors influencing the strength and the number of Lewis acid sites to be distinguished: (a) A cation effect: comparing spectra of adsorbed pyridine on Cr, Fe, Al, Ga HTB compounds, we observe that the strength of the sites varies in the order Al > Ga > Cr Fe. This order is logical considering the w/r2 ratio, where w is the

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Functionalized Inorganic Fluorides

electronegativity on the Allred-Rochow scale and r is the ionic radius of the cation. The same order is observed for the thermal stability of the compounds.51 (b) Effect of the amount of fluorine: the frequencies of the 8a and 19b bands for coordinated pyridine are at higher wavenumbers than those observed on the corresponding oxides. It is well known that the electron attracting effect of fluorine increases the acid site strength. This becomes very important, as revealed by the high positions of the 8a and 19b pyridine species at 1628 and 1457 cm1, respectively, approaching the value observed in the py,BF3 complex. This point also explains why the hydroxy fluorides are less acidic than their corresponding fluorides. In particular, the strength of pyrochlore compound (F/Al ¼ 1.8) acid sites is less than those of the HTB compounds (F/Al ¼ 2.6) but greater than those of the corresponding oxides.49 (c) Effect of the structure: when the pores of the material are sufficiently large (as for example in the FeIII HTB compound), ammonia adsorption shows the presence of Lewis acid sites inside the material.52 This increases considerably the number of sites and demonstrates the presence of poorly accessible anionic vacancies. The number of Lewis acid sites at the surface of a range of aluminium fluoride derivatives (quantified through the intensity of the 19b band of strongly coordinated pyridine) has been shown to be related directly to the specific surface area of the compound.48 This finding is particularly relevant when the effectiveness of different Lewis acids, say as catalysts, is being compared and it illustrates very well the importance of the synthetic method used to produce high surface area materials. In many compounds the Lewis acid sites are highly heterogeneous. The use of CO as probe molecule has shown the presence, in addition to the Brønsted type discussed above, of five types of Lewis acid sites on the surface of the AlF2.6(OH)0.4 high surface area HTB compound (Figure 4.13).54 They are labelled A (2235 cm1, D ¼ 94 cm1), B (2215 cm1, D ¼ 74 cm1), C (2200 cm1, D ¼ 59 cm1), D (2183 cm1, D ¼ 42 cm1) and F (2166 cm1, D ¼ 25 cm1). Postfluorination of the sample using CHF3 modifies the numbers of some types of sites; sharp decreases of A, C, D and E sites were observed, while the quantity of B sites was affected to a far smaller extent. To help in the assignment of CO adspecies, a theoretical modeling investigation has been performed. From the morphology of an HTB b-AlF3 crystallite it can be predicted that the exposition of (010) and (001) surfaces is preferential, whereas the (100) surface is unstable.56 Ab initio calculations simulating CO adsorption on the partially hydroxylated, (010) face show that the wavenumber of the coordinated CO species is the highest when the coordinatively unsaturated AlIII pentacoordinated site is surrounded only by fluorine anions. The substitution of F ions by OH groups clearly results in a shift to lower wavenumber. As a consequence, the B site in Figure 4.13 should correspond to AlF5 entities, the C site to a AlF4(OH) environment, while the sites D and F are due to AlF3(OH)2 environments. These assignments agree well with the modifications in band intensities after fluorination. To complete the surface description, the A site will probably correspond to very unstable AlF5 entities on the (100) face, or to Al3þ ions in a tetrahedral coordination. As has been described above, the heterogeneity of the Lewis sites essentially arises from the partially hydroxylated state of many of the aluminum fluorides. A greater

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homogeneity of the surface can be obtained via a postfluorination treatment, as a sample thermal treatment under CF3H. The number of strong acid sites so obtained depends on the probe molecule used to measure them (CO: 0.2–0.4 site per nm2; py: 1.2–1.5 site per nm2).45, 48, 53 These results for HTB-AlF2.6(OH)0.4 will be reconsidered below, in the light of complementary radiotracer experiments, to discuss in a complete way the surface behaviour of this fluoride using a variety of methods.

4.4.2 4.4.2.1

Using HCl as a Probe with Detection via [36Cl]-Labelling Aluminium Chlorofluoride

The benchmark compound for this technique is aluminium chlorofluoride (ACF), some of whose properties were described in Section 4.2.2 above. There are many pieces of evidence for this compound being a strong, solid Lewis acid.19, 23 Exposure of ACF to successive aliquots either of anhydrous H36Cl or of [36Cl]-ButCl leads to rather similar behaviour, shown in Figure 4.14. This describes the [36Cl] surface counts resulting from H36Cl (a) or [36Cl]-ButCl additions (b). Although numerical comparisons between (a) and (b) are not possible, since the [36Cl] specific activities of the H36Cl and [36Cl]-ButCl used differ, the very similar qualitative relationship observed suggests that the surface count data observed in (b) are due to adsorbed H36Cl, which results directly from dehydrochlorination of [36Cl]-ButCl at the ACF surface. Removal of volatile material by condensation after count No. 9 in Figure 4.14 (a), leads to a reduction in the surface count, No. 10, and a further reduction after one day, count No. 11. However, the count is still well above background, indicating that a significant fraction of H36Cl is strongly adsorbed. A small increase in the [36Cl] surface count derived from exposures of [36Cl]-ButCl to ACF, count No. 9 in Figure 4.14 (b), is observed after removal of all volatiles, i. e. after count No. 8. This can be accounted for by the assumption that some dehydrochlorination sites with the accompanying adsorption sites for H36Cl are located at grain boundaries within the solid ACF rather than being limited to the exterior surface. The isotope [36Cl], in common with all b emitters, is subject to self-absorption of its radiation.42 It will be detected by the endwindow G. M. counters used (Figure 4.10) only when located at the exterior surface of the sample under investigation. Evidently in the case of ACF, some time must elapse for migration of the [36Cl] species to occur from the bulk to the exterior surface. This phenomenon is observed also for the other compounds described in this section. 4.4.2.2

HS-Aluminium Trifluoride

As explained above HCl can be used directly or as a product from the dehydrochlorination of ButCl. The latter reaction is easily observed over HS-AlF3 at room temperature57 and comparisons between the behaviour of HS-AlF3 with those of fluorinated (by SF4) -alumina and b-AlF3 (HTB structure) can be made.41,44 They indicate that the greatest reactivity is found for HS-AlF3, either because it has the greatest site density (specific surface areas of samples of this compound as determined by the BET method are usually in the range 200–300 m2 g1)45 or because it has uniquely strong Lewis acid sites. The

124

Functionalized Inorganic Fluorides 10 500 (a)

10 000

500 s surface count

9500 9000 8500 8000 7500 7000 6500 0

2

4

6

8

10

12

count No. 1–9 (1.3 kPa); 10, 11 (after gas removal) 25 000

500 s surface count

24 000

23 000

(b)

22 000

21 000

20 000

0

2

4

6

8

10

count No.1–8(6.7 kPa); 9 (after vapour removal)

Figure 4.14 [36Cl] Surface count relationships from the successive additions of (a) H36Cl and (b) [36Cl]-ButCl to aluminium chlorofluoride (ACF). Line breaks correspond to the removal of the last aliquot of vapour. (Reprinted with permission from M. Nickkho-Amiry and J. M. Winfield, J. Fluorine Chem., 128, 344–352 Copyright (2007) Elsevier Ltd.)

behaviour of H36Cl, either from direct addition or generated in situ, indicates both physical and chemical adsorption.57 In the latter case both exterior (at the gas-solid interface) and interior (boundaries between particles) surfaces are involved. This distinction is possible experimentally due to the self-absorption property of the b radiation emitted from [36Cl] as discussed above for ACF. The exact behaviour observed when [36Cl] surface counts are examined over a range of H36Cl additions depends on the fluorinating agent used to prepare HS-AlF3 from the initial sol-gel. The behaviour of HS-AlF3 samples, which have been prepared either by fluorination with CCl2F2/N2 or by fluorination with aHF/N2 in the second stage of their preparations, towards H36Cl additions is compared in Figure 4.15 (a) and (b).

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Exposure of HS-AlF3, which has been prepared by the CCl2F2/N2 route, to successive aliquots of H36Cl, Figure 4.15 (a), leads to a progressive build up of [36Cl] on the surface. Its removal, in vacuo with occasional pumping, is slow enough to be monitored; the initial loss, count Nos. 9–11 is relatively rapid but even after five days, count No. 12, a substantial fraction remains. Evidently much of the H36Cl is strongly bound. The BET area of the material synthesized via the aHF/N2 route is larger, reflected in the larger surface counts shown in Figure 4.15 (b). In this case the [36Cl] surface counts are almost constant over the range of H36Cl additions; the value, count No. 10, measured immediately after removal of the last aliquot, falls to a low level but counts made 24 h later, count Nos 11 and 12, are 10 500

(a) 10 000

500 s surface count

9500 9000 8500 8000 7500 7000

0

28 000

2 4 6 8 10 12 count No. 1–3 (6.0 kPa), 4–8 (8.0 kPa), 9–12 (during gas removal)

(b)

26 000

500 s surface count

24 000 22 000 20 000 18 000 16 000 14 000 12 000 10 000 8000 0

2 4 6 8 10 count No. 1–9 (5.3 kPa), 10–12 (after gas removal)

12

Figure 4.15 [36Cl] Surface counts from an high surface area (HS) AlF3 sample, whose precursor was fluorinated with (a) CCl2F2/N2, or (b) with HF/N2, after exposure to sequential aliquots of H36Cl. Line breaks correspond to the removal of the last aliquot of gas. (Reprinted with permission from M. Nickkho-Amiry, G. Eltanany, S. Wuttke, S. Ru¨diger, E. Kemnitz and J.M. Winfield, J. Fluorine Chem., 129, 366–375 Copyright (2008) Elsevier Ltd.)

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greater. This appears to reflect diffusion of [36Cl] species from bulk to the exterior surface, as proposed above for ACF. A more extensive comparison of the two types of HS-AlF3 is provided by the dehydrochlorination experiments shown in Figure 4.16 (a) and (b).

(a)

10 000

500 s surface count

9000 8000 7000 6000 5000 –2 0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 count No. 1–8, 10–17 and 19–26 (all 5.3 kPa); 9, 18 and 27 (vapour removed)

18 000 16 000 500 s surface count

14 000 12 000

(b)

10 000 8000 6000 4000 2000 0

–2 0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 count No. 1–8, 11–19, 21–29 (5.3, 2.7, 8.0 kPa); 9–10, 20, 30–31(after vap. removed)

Figure 4.16 Comparison of the behaviour of [36Cl]-ButCl towards HS-AlF3 samples whose precursors had been fluorinated with (a) CCl2F2/N2 and (b) with HF/N2; three consecutive sets of sequential exposures are shown for each sample. Line breaks correspond to the removal of the last aliquot of vapour in a set and before commencing the next. (Reprinted with permission from M. Nickkho-Amiry, G. Eltanany, S. Wuttke, S. Ru¨diger, E. Kemnitz and J.M. Winfield, J. Fluorine Chem., 129, 366–375 Copyright (2008) Elsevier Ltd.)

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127

In both cases the [36Cl] surface count data are the result of a series of three series of experiments in which eight or nine aliquots of [36Cl]-ButCl were allowed successively to contact a sample of HS-AlF3.57 After each series of additions, the [36Cl] count from the surface in the absence of vapour was recorded, count Nos 9, 18 and 27 in Figure 4.16 (a) and 9–10, 20 and 30–31 in Figure 4.16 (b). Although in neither case is there any substantial evidence for the inhibition of dehydrochlorination by strongly bound H36Cl, it appears that retention of strongly bound H36Cl is relatively less important by HS-AlF3 prepared via the aHF route, Figure 4.16 (b). The difference between the two samples of HS-AlF3 can be rationalized in terms of HF-blocking of the strongest Lewis sites at which, in the absence of HF, HCl would be adsorbed. It appears also that HCl/HF adsorption sites are distinct from sites at which dehydrochlorination occurs.

4.4.2.3

HS-Magnesium Difluoride and Related Materials

Magnesium fluoride is usually regarded as a Lewis base fluoride because in its reactions it is a potential source of F anion. As conventionally prepared, it is not known to be a Lewis acid. It has a moderate specific surface area, 45 m2 g1 and has been used as a support material for heterogeneous catalysts; in this latter context, surface hydroxyl groups generated during manipulation of MgF2 can behave as Brønsted bases.58 However, as prepared by the sol-gel route,46 its behaviour towards H36Cl and [36Cl]-ButCl is rather similar to that of HS-AlF3 as described above. However, it is less active with respect to room temperature dehydrochlorination. For example there is no evidence from FTIR examination of the vapour above HS-MgF2 for HCl formation after ButCl has been admitted, although it is detected easily at the surface by [36Cl] labelling.57 Similar behaviour is shown by the amorphous solid, HS-MgF2/15 mol% FeF3 (equivalent in a crystalline context to a Tanabe solid40,59 in which Lewis acidity is enhanced compared with MF2 by doping with M0 F3). Because of their similar BET areas and pore volumes, [36Cl]-surface counts on 15%FeF3 in HS-MgF2 can be compared with HS-MgF2; the comparison indicates that the 15% FeF3 in HS-MgF2 material is the more effective Lewis acid. A similar finding is indicated from NH3 TPD experiments.57 The behaviour of these two materials with respect to their behaviour towards H36Cl and [36Cl]-ButCl are compared in Figures 4.17 and 4.18.

4.4.2.4

Common Features

The fluoride surfaces whose behaviour has been described above all exhibit the following features when examined by the Geiger-Mu¨ller direct monitoring method: (a) Surface radioactivity is readily detected at each surface following either direct exposure of H36Cl or [36Cl]-ButCl to the thinly spread solid at room temperature. (b) The counts are detectable in the [36Cl]-ButCl case even when the evolution of HCl to the gas phase above the solid is too small to be detected by transmission FTIR. This illustrates the sensitivity of the method but it is aided also by the relatively large BET areas that these solids have.

128

Functionalized Inorganic Fluorides

(c) The [38Cl]-surface activity deposited at the surface is, to a great extent, strongly bound. It is located also within the ‘bulk solid’; this location is not detected by the counters due to self-absorption of the b emitter, [36Cl] but its existence is inferred by the surface count behaviour following removal of radioactive vapour and weakly bound [36Cl] species. Possible ways of formulating the adsorbed states of H36Cl are discussed below (in Section 4.4.4). 26 000 (a)

500 s surface count

24 000 22 000 20 000 18 000 16 000 14 000 12 000 10 000 0

2 4 6 8 10 count No. 1–8 (3.3 kPa), 9,10 after gas removal

95 000

500 s surface count

90 000 (b) 85 000 80 000 75 000 70 000 0

2 4 6 8 10 count No. 1–8 (2.7 kPa), 9,10 (after gas removal)

Figure 4.17 Comparisons among [36Cl] surface counts from H36Cl aliquots in contact with (a) high surface area (HS) MgF2 and (b) 15 mol% HS-FeF3 in HS-MgF2. Line breaks correspond to the removal of the last aliquot of gas. (Reprinted with permission from M. Nickkho-Amiry, G. Eltanany, S. Wuttke, S. Ru¨diger, E. Kemnitz and J.M. Winfield, J. Fluorine Chem., 129, 366–375 Copyright (2008) Elsevier Ltd.)

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129

7000 (a) 500 s surface count

6500 6000 5500 5000 4500 4000 3500 0

2

4

6

8

10

count 1–8 (4.0 kPa), 9,10 (after vapour removal)

14 000 (b)

500 s surface count

13 000 12 000 11 000 10 000 9000 8000 7000 0

2

4

6

8

10

count No. 1–8 (4.0 kPa), 9 (after vapour removed)

Figure 4.18 Comparisons between [36Cl] surface counts from [36Cl]-ButCl aliquots in contact with (a) HS-MgF2 and (b) 15 mol% HS-FeF3 in HS-MgF2. Line breaks correspond to the removal of the last aliquot of vapour. (Reprinted with permission from M. Nickkho-Amiry, G. Eltanany, S. Wuttke, S. Ru¨diger, E. Kemnitz and J.M. Winfield, J. Fluorine Chem., 129, 366–375 Copyright (2008) Elsevier Ltd.)

4.4.3

Metal Fluoride Surfaces that Contain Surface Hydroxyl Groups: Aluminium Hydroxy Fluorides with the Hexagonal Tungsten Bronze Structure

The focus in this section is on derivatives of the metastable b phase of aluminium trifluoride, which has the hexagonal tungsten bronze structure with AlF6 octahedra linked

Functionalized Inorganic Fluorides

130

to form hexagonal hollow tubes running through the lattice50 (see also Figure 4.12). In addition to its use as a laboratory heterogeneous catalyst,1 it can be considered to be a relative of amorphous HS-AlF3; this is illustrated, for example, by a recent comparative study of vibrational spectra of AlF3 forms.60 b-Aluminium trifluoride was one of the earliest fluoride Lewis acid surfaces to be investigated by the use of py as a probe species for surface acidity.61 Because of the way in which it is prepared, b-AlF3 usually has water located in its hexagonal channels.50 Water can migrate to the surface where it can coordinate to Lewis acid AlIII sites or initiate hydrolysis to form hydroxylated AlIII sites. Both types of site can be probed using H36Cl;41,62 the steps that are believed to be involved are shown schematically in Figure 4.19. Al

Al F

F H2O

Al F

OH2

723 K/He

F Al

F

F

F Al

and

Al

then 523 K in vacuo

F

OH

Al

F

major

F minor

hydrated site H

+ HF

Al

hydroxylated site

H room temp.

O

+ H2O + HCl or pump

Al

Al F

>373 K

OH2.ClH

+ HCl

F

F

Al F

F

F

HCl at room temperature

Cl

OH room temp. + HCl

Al F

Al

to 623 K

F

F

+ H2O F

chlorination of hydroxyl sites

Figure 4.19 Schematic describing the proposed formation of hydrated and hydroxylated surface sites at HTB b-AlF3 and the subsequent adsorption/desorption of anhydrous HCl. (Reprinted with permission from C. H. Barclay, H. Bozorgzadeh, E. Kemnitz, M. NickkhoAmiry, D. E. M. Ross, T. Skapin, J. Thomson, G. Webb and J. M. Winfield, J. Chem. Soc., Dalton Trans., 40–47 Copyright (2002) Royal Society of Chemistry.)

It follows from these proposals that water coordinated to a surface AlIII site can give rise to Brønsted activity and on exposure to HCl behaves as a Brønsted base. The replacement of surface –OH groups by Cl, which is a result from HCl treatment above room temperature,62 is reminiscent of the chlorination of alumina by HCl38, 63 (cf. Figures 4.6 and 4.7).

Investigation of Surface Acidity

131

A more detailed examination becomes possible for compounds where hydroxyl groups are deliberately incorporated into the structure as opposed to arising from trace hydrolysis. Here we remind the reader of several points that were made above in Section 4.4.1. In the HTB solids of stoichiometry, MF3x(OH)x, where M is Al, Cr or Fe,51 both Brønsted and Lewis acidity can be demonstrated by FTIR spectroscopy using the basic probe molecules, NH3 and py.52 Even more information is available if the synthesis route is modified to produce solids having greater surface areas, for example by use of the microwave-activated, solventothermal process.47 The HTB-structured solid whose stoichiometry is AlF2.6(OH)0.4 is a particularly good example and its surface acidity has been studied using pyridine, lutidine (2,6-dimethylpridine), CO, H36Cl and [36Cl]-ButCl as surface probes.54 The pertinent FTIR spectroscopic data for the b-AlF2.6(OH)0.4 surface, which were discussed in Section 4.4.1, are summarised for convenience in Table 4.2. Making comparisons with related studies on aluminas,6, 9, 64 the presence of strong surface Lewis acidity is evident and there is good evidence for Brønsted acidity associated with the –OH groups.

Table 4.2 Characterization by FTIR spectroscopy of the Lewis/Brønsted acidic sites displayed by the b-AlF2.6(OH)0.4 surface. Data taken from reference 54. (Reproduced by permission of the PCCP Owner Societies.) Probe IR bands

Lutidine 8a (CC)/cm1 Pyridine 8a/cm1 CO (CO)/cm1

Bands due to Lewis acidity

1615–1620 1610

1628 1620

Bands due to Brønsted acidity

1652, 1631a

–

Proton affinity/kJ mol1

963

a

912

2235 2220–2215 2200 2183 2166 2173 D(OH) 160 598

Band is v8b.

The use of several probe molecules has enabled complementary information to be obtained, the detail of which depends on the identity of the probe. In particular, heterogeneity in the Lewis site strength is indicated on the exterior surface, as demonstrated by the frequency range and the basicity displayed by a probe molecule.54 Using py as the probe, two types of Lewis acid sites are detected: Figure 4.20; the first one, denoted L1 in Figure 4.20, is very strong ( 8a py band at 1628 cm1) but is present only in low concentration on the surface (approximately 0.2 sites nm2). The second type, denoted L2, is weaker ( 8a py band at 1620–1623 cm1) but is more abundant (approximately 1.2 sites nm2). Adsorption of CO, which was illustrated in Figure 4.13, indicates the existence of five different types of Lewis site (A-D and F) and a Brønsted site (E), the latter being responsible for the shift induced in the (OH) band (Figure 4.13, inset A). Pyridine-CO co-adsorption experiments reveal the relationships among the two sets of individual

Functionalized Inorganic Fluorides

132

A

L2

3680

B

Abs 0.01 L1 (c-b)

0.05 3500

3000

2500

/cm–1 1635

1625

1615 /cm–1

3680 1454

0.5

(a)

Abs 0.04

1620

1600 (b) (c) (d) (e) (f) 1628

1457

(g) 3680

3640 /cm–1

3600

1650

1600

1550 /cm–1

1500

1450

Figure 4.20 IR spectra of b-AlF2.6(OH)0.4 after activation at 573K before (a) and after introduction of an equilibrium pressure (133 Pa) of pyridine (b-g); (b) evacuation at room temperature under vacuum and thermodesorption at (c) 323, (d) 423, (e) 473, (f) 523 and (g) 573 K. Inset A: difference IR spectra after introduction of an equilibrium pressure (133 Pa) of pyridine followed by evacuation at r.t. Inset B: deconvolution (dotted lines) of the n8a vibrational mode (at room temperature). (Reprinted with permission from D. Dambournet, H. Leclerc, A. Vimont, J.-C. Lavalley, M. Nickkho-Amiry, M. Daturi and J. M. Winfield, Phys. Chem. Chem. Phys. 11, 1369–1379 Copyright (2009) PCCP Owner Societies.)

experiments. The L1 and L2 types identified from py adsorption/desorption, Figure 4.20, are both heterogeneous; L1 and L2 sites are each related to two sites. Site A and site B, identified by (CO) bands at 2235 and 2215 cm1, are included in the L1 envelope; site C and site D, characterized by (CO) bands at 2200 and 2183 cm1 are related to L2. Adsorption of CO reveals also the presence of weaker Lewis acid sites characterized by a (CO) band at 2166 cm1 (site F in Figure 4.13). The different types, and strengths, of the Lewis sites are suggested to arise as a consequence of the various F/OH anionic environments that exist, presumed to be related to those revealed previously in the bulk by 27Al NMR.47 The IR spectrum of lutidine (DMP) adsorbed on b-AlF2.6(OH)0.4 and displayed in Figure 4.21, shows the presence of coordinated lutidine ( 8a at 1615, 8b at 1590 cm1) and lutidinium species ( 8a at 1652 and 8b at 1630 cm1). The latter bands indicate that surface Brønsted acid sites are present. In agreement with this

Investigation of Surface Acidity

133

conclusion, subtracted spectra in the hydroxyl stretching region, suffer a decrease in the intensity of the 3680 cm1 band (Figure 4.21, inset). Despite the steric hindrance induced by the methyl groups of lutidine, it appears that accessible acidic OH groups are present at the surface. Abs 0.2 (a) (b)

3680

(b)-(a) 3680

3640 /cm–1

3600

Physisorbed Abs 0.1

Lewis Brønsted 1615 1652 1631 Adsorption Evacuation RT 1620

373 K

423 K 1700

1600 1650 Wavenumber/cm–1

1550

Figure 4.21 Difference IR spectra of coordinated lutidine on b-AlF2.6(OH)0.4 after adsorption of 133 Pa at equilibrium pressure, followed by desorption under vacuum at room temperature (293 K), 373 and 423 K. The dotted curve is the difference curve between desorption obtained at 373 and 473 K. Inset: spectra in the n(OH) region of the sample activated at 573 K (a) and after adsorption of lutidine (b). (Reprinted with permission from D. Dambournet, H. Leclerc, A. Vimont, J.-C. Lavalley, M. Nickkho-Amiry, M. Daturi and J. M. Winfield, Phys. Chem. Chem. Phys. 11, 1369–1379 Copyright (2009) PCCP Owner Societies.)

To summarize, experiments with CO and lutidine, described above, both indicate that surface OH groups ((OH) at 3680 cm1) are present, which generate homogeneous Brønsted acid sites with a medium acid strength (D(OH) ¼ 160 cm1, (CO) ¼ 2173 cm1), able to protonate lutidine but not pyridine.

134

Functionalized Inorganic Fluorides

The behaviour of b-AlF2.6(OH)0.4 towards H36Cl is shown in Figure 4.22.54 From this sample, loosely bound H2O, including as much as possible of the water originally located in the hexagonal channels, had been removed before exposure to H36Cl. The interaction of H36Cl is progressive with increasing number of aliquots used and both surface and bulk appear to be involved. Most obviously, the interaction should involve Al-OH groups both those located at the exterior surface and in the hexagonal channels. Migration of H36Cl appears to occur during the storage period in vacuo, (from count No. 9 to count No. 10 in Figure 4.22) from hexagonal channels in the HTB-structure, where [36Cl] b radiation cannot be detected because of self absorption42 of [36Cl], to sites on the exterior surface where it is detectable.

13 000

500 s surf. count

12 000 11 000 10 000 9000 8000 7000 6000 0

2 4 6 8 10 count No. 1–8 (3.3 kPa), 9,10 after gas removal

Figure 4.22 Surface counts from anhydrous b-AlF2.6(OH)0.4 during exposure to H36Cl. The line break corresponds to the removal of gas. (Reprinted with permission from D. Dambournet, H. Leclerc, A. Vimont, J.-C. Lavalley, M. Nickkho-Amiry, M. Daturi and J. M. Winfield, Phys. Chem. Chem. Phys. 11, 1369–1379 Copyright (2009) PCCP Owner Societies.)

The activity of b-AlF2.6(OH)0.4 with respect to room temperature dehydrochlorination of [36Cl]-ButCl is significantly smaller than that observed for HS-AlF3 or even, possibly, for HS-MgF2. Deposition of [36Cl]-species on the surface, although observed, could not be quantified.54 However, deposition of H36Cl could be achieved using a sample of b-AlF2.6(OH)0.4 from which water had not been exhaustively removed, b-AlF2.6(OH)0.4,xH2O, x is approximately 0.12. It appears therefore that some Lewis acid sites are present even in the presence of the base H2O. The apparent promotional role of H2O in the room temperature dehydrochlorination of ButCl may be the result of migration of the HCl liberated to a nearby water molecule, where it will be strongly adsorbed (cf. Figure 4.19). The presence of hydroxyl groups introduces an extra dimension to the chemistry as OH groups, both located at the surface and in the channels, can behave as Brønsted acids,

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135

towards CO or lutidine and as Brønsted bases towards the potentially strong Brønsted acid, HCl. Although less reactive with respect to dehydrochlorination than, for example HS-AlF3, retention of H36Cl is substantial. The obvious comparison is with the behaviour of H36Cl at conventionally prepared b-AlF3, where the slow desorption of H2O.HCl has been observed at room temperature (Figure 4.19).41 The behaviour of the HTB hydroxy fluorides also provides a bridge to the behaviour of partially hydrolysed metal fluoride surfaces, to be discussed below.

4.4.4

Possible Geometries for HCl Adsorbed at Metal Fluoride Surfaces: Relation to Oxide and Oxyfluoride Surfaces

Throughout this account of surface acidity we have attempted to emphasize the relationships among oxides and fluorides in the context of acidic surfaces. The development of high surface-area metal fluorides makes this comparison far easier than has previously been possible. The various types of interaction between such surfaces and anhydrous HCl are cases in point. This aspect was introduced in Section 4.3.2.2 above and possible modes of adsorption of HCl at metal fluoride surfaces were sketched in Figure 4.8. Anhydrous HCl is commonly dissociatively adsorbed at aluminas. For example, analysis of vibrational data for HCl adsorbed at -alumina63 led to the pictorial representations of HCl adsorbed at medium strong and at a site approximating to a weak Lewis acid type (cf. Figure 4.1) shown in Figures 4.6 and 4.7 respectively. In addition there is vibrational spectroscopic evidence for associative adsorption, molecular HCl being either physisorbed or weakly chemisorbed; in either event the species is desorbed completely by 423 K.63 Dissociative adsorption is also a likely outcome when HCl is exposed to high surface area oxyfluorides. This has been studied at room temperature using H36Cl and the direct G. M. monitoring technique (Figure 4.10) with a series of fluorinated chromia and lightly-doped MgII and ZnII chromia aerogels.65 Significant ( 60–97%) fractions of the H36Cl aliquots were retained by the surfaces, the behaviour being comparable to that observed towards -alumina. Associative adsorption is a possibility at the surfaces of binary fluorides such as HS-AlF3 (Section 4.4.2.2) and HS-MgF2 and its derivatives (Section 4.4.2.3). These situations have not yet been studied by vibrational spectroscopy, so the inference is made solely on the basis of [36Cl] experiments. Associative adsorption requires HCl to behave as a Lewis base and this type of behaviour is expected only if very strong Lewis acid sites are present.57 However, there are two other possible explanations for the significant degrees of H36Cl retention that are observed. Firstly, it is possible to envisage hydrogen bonding between HCl and surface fluoride and secondly, and perhaps more plausibly, HCl could interact with surface hydrated or hydroxylated sites that are the result of hydrolysis to give rise to a dissociatively adsorbed species. This situation is envisaged to be relevant to the HTB solids b-AlF2.6(OH)0.4 and b-AlF2.6(OH)0.4.xH2O, as some of the Lewis acid sites in b-AlF2.6(OH)0.4 will have mixed F/OH nearest neighbours.54 The possible adsorbed states for HCl at a hypothetical fluoride/hydroxide surface, where the solid is comprised of an aggregate of nanoparticles, is shown diagrammatically in Figure 4.23. Although the picture is speculative, the composite is consistent with all the observations on these solids made to date.

136

Functionalized Inorganic Fluorides F Al

Al F

OH Al

F

F

Al O

F

HCl

ClH

F Al

Al HCl

OH2 Al F

F

HCl

Cl

HCl

Al O

F

HCl exterior surface

HCl

HCl

interior surface HCl

Figure 4.23 Schematic of HCl adsorption at an hypothetical fluoride-hydroxide surface with the bulk having an aggregated nanoparticle structure

4.5

Conclusions

In this chapter recent developments in the study of the acidity in metal fluorides have been surveyed. The relatively large specific surface areas encountered aid the study of acidic metal fluorides both by FTIR spectroscopy and by a reaction-specific radiotracer approach. Where the structure of the solid is known, the enhanced surface area and the application of both techniques enables a relatively detailed picture of the acidic properties to be obtained. This is the situation for the HTB-aluminium hydroxy fluoride, Figure 4.24. FTIR studies with a range of Lewis base probe molecules point to the presence of a range of Lewis acid sites in HS-metal fluorides, some of which are of higher acidity than in the corresponding oxides. The application of a radiotracer method in which [36Cl] labelled species are used to track species formed at the surface, is less direct than the FTIR method, as structural information is not possible. However this approach does enable features of the bulk to be related to surface events. Because it is also amenable to the use of vacuum and dry box techniques, it is very useful for the study of hydrolytically sensitive fluoride surfaces. In many respects therefore, the radiotracer and FTIR approaches to the examination of surface acidity are complementary. Information that relates to chemical speciation cannot

Investigation of Surface Acidity

= F/OH Al

137

Me

Al

N

Me Me

Cl Al

Al

ClH

N

Al

Al

CO

Figure 4.24 A schematic of the investigation of HTB AlF2.6(OH)0.4 using adsorption of the probe molecules CO, pyridine, lutidine and [36Cl]-hydrogen chloride. (Reprinted with permission from D. Dambournet, H. Leclerc, A. Vimont, J.-C. Lavalley, M. Nickkho-Amiry, M. Daturi and J. M. Winfield, Phys. Chem. Chem. Phys. 11, 1369–1379 Copyright (2009) PCCP Owner Societies.)

be obtained directly from radiochemical monitoring but is directly available by FTIR. The use of a b emitter enables the fate of an adsorbed species to be ‘tracked’ in favourable circumstances from the bulk to the surface.

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5 Probing Short and Medium Range Order in Al-based Fluorides using High Resolution Solid State Nuclear Magnetic Resonance and Parameter Modelling Christophe Legein1, Monique Body2, Jean-Yves Buzare´2, Charlotte Martineau1,3 and Gilles Silly4 1

Laboratoire des Oxydes et Fluorures, CNRS UMR 6010, Institut de Recherche en Inge´nierie Mole´culaire et Mate´riaux Fonctionnels, CNRS FR 2575,Universite´ du Maine, Avenue Olivier Messiaen, 72085 Le Mans Cedex 9, France 2 Laboratoire de Physique de l’Etat Condense´, CNRS UMR 6087, Institut de Recherche en Inge´nierie Mole´culaire et Mate´riaux Fonctionnels, CNRS FR 2575, Universite´ du Maine, Avenue Olivier Messiaen, 72085 Le Mans Cedex 9, France 3 Tectospin, Institut Lavoisier de Versailles (UMR 8180), Universite´ de Versailles Saint Quentin en Yvelines, 45 Avenue des Etats-Unis, 78035 Versailles cedex, France 4 Institut Charles Gerhardt Montpellier, UMR 5253, CNRS-UM2-ENSCM-UM1, Physicochimie des Mate´riaux De´sordonne´s et Poreux, Universite´ de Montpellier II, Place Euge`ne Bataillon, C.C. 1503, 34095 Montpellier Cedex 5, France

5.1

Introduction

We show in this review that advanced solid state nuclear magnetic resonance (NMR) methods coupled with empirical and ab initio calculations of relevant parameters are well placed to meet the challenges provided by modern material chemistry. We demonstrate the

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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ability of these techniques on Al-based fluorides not only to characterize ordered crystalline structures but also to give detailed insights into the structure of partially disordered, amorphous or glassy materials. For these compounds, solid state NMR is a privileged technique since they contain at least two 100 % abundant and sensitive nuclei, 19F (nuclear spin I ¼ 1/2) and 27Al (I ¼ 5/2) which is a quadrupolar nucleus. This chapter is divided into three parts. In the first part (Section 5.2), various solid state NMR techniques, such as magic angle spinning (MAS), cross polarization (CP), multiple-quantum (MQ) and multidimensional spectroscopies, pertinent to obtain high resolution NMR spectra in Al-based fluorides, are presented. They allow the precise determination of the 19F isotropic chemical shift iso and 27 Al iso and quadrupolar parameters or reveal 19F-19F or 19F-27Al proximities. In the second part (Section 5.3), the strong potential of these techniques for the study of functionalized Al-based inorganic fluorides, synthesized for their possible application as catalysts (fluorinated alumina and zeolites, high surface area aluminium trifluorides and nanostructured aluminium fluoride hydrates and hydroxyfluorides), is shown. The compounds used as model for the assignment of the 19F and 27Al NMR lines in these materials are also presented. The third part (Section 5.4) is devoted to the determination of the 19F and 27Al NMR parameters of crystalline alkali and alkaline-earth fluoroaluminates. In the crystalline phases, the modelling of these parameters, starting from diffraction data and using semiempirical and Density Functional Theory (DFT) methods, aids spectral assignment. For 19F, double-quantum single-quantum (DQ-SQ) MAS NMR correlation experiments are demonstrated to be a complementary tool for correct line assignments.

5.2

High Resolution NMR Techniques

The quest for increased spectral resolution is a common feature of many solid-state NMR spectroscopists and has induced the developments of higher homogeneous magnetic fields (22 T spectrometers are now commercially available) and faster MAS (now up to 70 kHz) and of new NMR experiments, the ultimate goal being to obtain a spectrum on which all individual spectral signatures are resolved. 5.2.1

Fast MAS and High Magnetic Field

High resolution 1D 19F MAS NMR spectra can usually be recorded at moderate magnetic fields (7–10 T) under fast MAS (30–35 kHz), combined with heteronuclear decoupling when necessary. Nevertheless, because of the high gyromagnetic ratio and 100 % natural abundance of the 19F nucleus, 19F MAS NMR spectra may suffer from homogeneous line broadening due to residual homonuclear dipolar couplings not averaged out by fast MAS. This is shown in Figure 5.1 for b-CaAlF5 [1]; this compound contains five inequivalent F sites and only four resonances are resolved on the 19F MAS NMR spectrum recorded at 7 T and 35 kHz. The use of higher magnetic fields and/or ultra fast MAS (up to 70 kHz) allows this limitation to be overcome. In b-CaAlF5 [1,2], the five 19F resonances are resolved at

Short and Medium Range Order in Al-based Fluorides

35 kHz

* **

* **

* **

30 kHz

* **

* * *

25 kHz

* **

* *

20 kHz

* **

*

* * * **

15 kHz

* **

*

* *

*

10 kHz 200

150

143

**

*

100

*

**

50

0

**

* *

* **

**

–50

* –100

* –150

Isotropic chemical shift (ppm) Figure 5.1 -CaAlF5 19F MAS NMR spectra collected at 7 T and at various spinning frequencies. The star symbols indicate the spinning sidebands [1]. Reprinted with permission from Inorg. Chem., 43, 2474–2485 (2004). Copyright 2004 American Chemical Society

17.6 T (Figure 5.2). For Na5Al3F14 [3] a dramatic improvement in resolution is also obtained at 19.6 T and very fast MAS (40 kHz). The increase in field enlarges the chemical shift difference (units: Hz) between resonances, reducing the spectral overlap between resonances. For quadrupolar nuclei, i.e. nuclei with spin > 1/2, an additional difficulty arises from the field-dependent shift and the line broadening by the second-order perturbation of quadrupolar interactions under the dominant Zeeman interaction [4]. MAS assists

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Functionalized Inorganic Fluorides

17.6 T

*

7T

35

30

25

20

15

10

5

0

–5

–10

Isotropic chemical shift (ppm)

Figure 5.2 19F MAS NMR spectra of -CaAlF5 collected at 7 T for r ¼ 35 kHz [1] and collected at 17.6 T for r ¼ 34 kHz [2]. The symbol * indicates a nonidentified impurity

resolution but does not completely remove the anisotropic broadening if the quadrupolar interaction is large [5]. The second-order quadrupolar effect being inversely proportional to the square of the Larmor frequency, a quadratic gain both in resolution and sensitivity is expected at higher magnetic fields [6]. This is of major importance in terms of applicability of solid-state NMR to crystalline and amorphous inorganic materials containing quadrupolar nuclei such as 27Al [7]. This is shown in Figure 5.3 for nanostructured aluminium

7T 17.6 T

40

0

–40 δiso (ppm)

–80

–120

Figure 5.3 Central transition of the 27Al NMR spectra of the low surface area (LSA) AlF2.4(OH)0.6 sample recorded at 7 T and 17.6 T. The spinning frequency is 30 kHz for both [8]. Reprinted with permission from Chem. Mater., 20, 1459–1469 (2008), copyright 2008 American Chemical Society

Short and Medium Range Order in Al-based Fluorides

145

hydroxyfluorides derived from b-AlF3 [8]. The high magnetic field, by reducing the influence of the contribution of the quadrupolar interaction to the spectrum, allows the resolution of 27Al resonances with close iso values. As shown in Section 5.3 of this chapter, 27Al high-field NMR is a powerful tool to probe and quantify the various aluminium species in aluminium based fluorides.

5.2.2 5.2.2.1

27

Al NMR SATRAS Experiments

Satellite Transition Spectroscopy (SATRAS) [9,10] has been proven to be a useful technique to determine the NMR parameters of quadrupolar spin systems, in particular when the quadrupolar frequency Q is so small that the structure of the central transition (CT) 1/2 ! 1/2 due to second-order effects is not resolved. In that case, reliable information can be obtained through the reconstruction of the spinning sideband manifold of the satellite transitions. The extent of the NMR spectrum gives Q while the shape of the envelope of the spinning sidebands provides the asymmetry parameter Q. When the effects of pulse duration can be neglected (i.e. using short pulse and small flip angle), the SATRAS spectrum is quantitative. As an example, the 27Al NMR spectrum of -BaCaAlF7, which contains one aluminium site, is presented in Figure 5.4. The shapeless central transition and the overall SATRAS spectrum expansion indicate a low Q value. The reconstruction of the SATRAS spectrum (Figure 5.4) was achieved using a homemade code based on the theoretical treatment developed by Skibsted et al. [9,11], and including a correction for the second-order frequency shift [12,13]. In this way, we precisely determined the 27Al NMR parameters (iso, Q and Q) in numerous alkali and alkaline-earth fluoraluminates [2,14–17].

5.2.2.2

MQ-MAS Experiments

In 1995, Frydman and Harwood [18] introduced the multiple-quantum MAS (MQ-MAS) experiment, which makes use of the multi-quanta transitions of a quadrupolar nucleus. This two-dimensional (2D) NMR experiment correlates a high resolution isotropic spectrum (F1 vertical dimension) to the anisotropic central transition of the MAS spectrum (F2 horizontal dimension). The development of this method has considerably extended the applicability of solid-state NMR to quadrupolar nuclei in multisite crystalline phases and glasses. In the presence of strong heteronuclear dipolar couplings (such as 19F-27Al), resolution of the 27Al MQ-MAS spectrum can also be improved by applying composite decoupling schemes during the MQ evolution and the acquisition periods, as shown for Na5Al3F14 [19]. In crystalline compounds, the separation of the aluminium sites in two dimensions allows extraction of their quadrupolar interaction parameters individually [2,14–17,19] (Figure 5.5). In disordered crystalline compounds, an improved resolution of the resonances is obtained, as shown on the 27Al MQ-MAS spectrum of a high surface area (HS) pyrochlore aluminium hydroxyfluoride [20].

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(b)

12

9

6

300

200

3

0

–3

–6

–9

–200

–300

–12

(a)

400

100 0 –100 frequency (kHz)

–400

Figure 5.4 Experimental (bottom) and reconstructed (top) spectra: (a) 27Al SATRAS NMR spectra [17] and (b) central transitions for -BaCaAlF7 collected at 7 T for r ¼ 10 kHz. In (a) the central transitions are truncated to improve the visibility of the spinning sidebands. (a) Eur. J. Inorg. Chem., 1980–1988 (2007), Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission

5.2.2.3

TOP Spectra

Another way to increase resolution while avoiding an excessively long experimental time is to construct a two-dimensional one pulse (TOP) spectrum [21–23] from a conventional 1D MAS spectrum. The TOP spectrum evidences the well-known narrowing of the inner satellite transition line width: when the components are not well resolved on the N ¼ 0 cross-section, the projection over the satellite transitions (N 6¼ 0) shows a much better resolution as shown in HS pyrochlore aluminium hydroxyfluoride [20] (Figure 5.6), nanostructured Al-based fluoride-oxide materials [24] and aluminium fluoride hydrate with cationic vacancies [25].

Short and Medium Range Order in Al-based Fluorides (a)

147

(b)

Isotropic dimension (kHz)

0 2 2

0

–2 –4 –6 –8 Frequency (kHz)

–10

–12

0

–2 –4 –6 –8 Frequency (kHz)

–10

–12

4

(c) 6 8

2

0

–2 –4 –6 –8 MAS dimension (kHz)

–10

–12 2

Figure 5.5 (a) Experimental 27Al 3Q-MAS NMR spectrum of g-BaAlF5 collected at (a) 7 T and r ¼ 25 kHz. The curves at the top and to the right of the spectrum are the full projections of the spectrum onto the MAS (F2) and isotropic (F1) dimension, respectively. The dashed lines indicate the F1 slices represented in (b) and (c). These slices are used for estimation of the NMR parameters. The reconstructed slices are shown below each experimental slice

3

(a)

4

(b)

2 1

N ≠ 0 ST

F1 (kHz)

12

16

1

20

2 3 4

24 20

Spinning sideband order

N = 0 CT

–15 –10 –5 0 5 10 15

10

0 –10 –20 –30 frequency (ppm)

27Al

–40

20

10

0 –10 –20 –30 frequency (ppm)

–40

27Al

Figure 5.6 (a) 3Q-MAS spectrum and (b) 2D TOP spectrum of HS pyrochlore aluminium hydroxyfluoride, AlF1.8(OH)1.2 0.3H2O. The N ¼ 0 cross-section shows the central transition spectrum and the N 6¼ 0 sum is the satellite transition spectrum showing enhanced resolution. The TOP spectrum was constructed by stacking subspectra shifted by the spinning frequency from a 1D MAS NMR spectra collected at r ¼ 30 kHz and 17.6 T [20]. J. Mater. Chem., 18, 2483–2492 (2008). Reproduced by permission of the Royal Society of Chemistry

148

Functionalized Inorganic Fluorides

5.2.3

High Resolution Correlation NMR Techniques

An important part of the power of solid-state NMR methods comes from their ability to produce homonuclear or heteronuclear correlation charts that can evidence spatial proximity from dipolar interaction [26,27] or chemical bonding from J-couplings [26,28]. On the one hand, the J-based NMR experiments make use of the nonvanishing isotropic part of the indirect J-coupling. Characterizing the existence of a chemical bond, these NMR methods can provide very detailed insight into the structure of crystalline or glassy materials [26,28–32]. However, to date, 19F-27Al J-coupling has never been observed or used in the solid state. On the other hand, the dipolar-based experiments make use of the dipolar interaction, which vanishes under MAS. Reintroduction of this through-space interaction (which depends inversely upon the distance between the spins to the third power) can be conveniently achieved using various recoupling scheme (continuous wave or modulated radio frequency pulses in CP experiments). 19F-27Al or 19F-19F dipolar-based experiments have been widely used to characterize Al-based fluorides and will be presented in the following.

5.2.3.1

1D CP-MAS Experiments

CP is used almost ubiquitously in NMR of spin-1/2 nuclei, principally as a means of enhancing signal to noise [33–34]. CP involving a quadrupolar nucleus is a more complex process, the dynamics of which are still poorly understood [5]. Although signal enhancements are rarely observed for many quadrupolar nuclei, the application of CP to quadrupolar systems still holds great potential as a spectral editing tool for the elucidation of spatial relationships. At the simplest level, the detection of a CP signal indicates a spatial proximity between two heteronuclei. Several 19F-27Al CP-MAS experiment results will be described in Section 5.3.

5.2.3.2

1D Double Resonance NMR Experiments

Among the double resonance experiments, REDOR (rotational echo double resonance) [35], TRAPDOR (transfer population in double resonance) [36,37] and REAPDOR (rotational echo adiabatic passage double resonance) [38] were applied to probe the presence and strength of heteronuclear 19F-27Al dipolar couplings in Al-based fluorides. The dipolar coupling between the nuclear spin species S (whose signal is detected) and spatially close nuclear spin species I is reintroduced into the experiments by coherent I-spin irradiation during the rotor period. If heteronuclear dipoledipole interactions are present, the I-spin irradiation causes a decrease in S-spin signal intensity, relative to a reference experiment without I-spin irradiation (intensity S0) as shown in Figure 5.7 [39]. The magnitude (S0 S)/S0 of the difference signal depends on both the strength of the dipole-dipole coupling and on the length of the overall dipolar evolution time. In an S{I} REDOR experiment dipolar recoupling is accomplished by applying p-pulses on the I channel in the middle of the rotor period. Although REDOR has successfully been applied to systems where one or other of the dipolar-coupled nuclei is quadrupolar (S or I),

Short and Medium Range Order in Al-based Fluorides

149

F-Ca(3)

Al-F-Ca(n)

a

control

difference 0

–50 –100 –150 relative frequency (ppm)

–200

–250

Figure 5.7 19F{27Al}TRAPDOR spectra of crystalline Ca2AlF7. Note the large change in relative intensities of the two groups of central peaks (filled in), representing F-Ca(3) and Al-F-Ca(n) fluoride sites. The label ‘a’ marks the signal from a small amount of crystalline CaF2 impurity, which also disappears in the difference spectrum [39]. Reprinted from J. Non-Cryst. Solids, 337, 142–149 (2004), copyright 2004 Royal Society of Chemistry, with permission from Elsevier

as for CP, the spin dynamics associated with the experiment become much more complex than those for the simple two-spin-1/2 case, and the experiment itself may become somewhat inefficient [5]. The REAPDOR and TRAPDOR experiments were designed specifically to study dipolar interactions involving quadrupolar nuclei. These sequences are similar to REDOR in their philosophy, preventing the refocusing of the dipolar interaction under MAS, but S{I} TRAPDOR experiments utilize continuous-wave irradiation and S{I} REAPDOR adiabatic mixing pulses on the I (quadrupolar) spin. 19 F{27Al} and 27Al{19F} REDOR experiments and 19F{27Al} TRAPDOR and REAPDOR experiments provided information on spatial proximities between 19F and 27 Al in many Al-based fluorides. We can cite aluminium fluoride phosphate glasses [40,41], aluminosilicate glasses [42] and zeolites dealuminated with NH4SiF6 [43,44]. Other examples are detailed in Section 5.3. To be exhaustive about 1D double resonance NMR experiments, we have to mention the TEDOR (transferred echo double resonance) experiment, which involves transferring magnetization between two heteronuclear spins [45]. Nevertheless, this experiment has been applied only once to Al-based fluoride, on AlPO4-CJ2, a microporous

150

Functionalized Inorganic Fluorides

oxy-fluorinated aluminophosphate, by Amoureux et al., as a preliminary experiment for designing the TEDOR-MQMAS method [46]. 5.2.3.3

2D Double Resonance NMR Experiments

In many cases, the spectral resolution achieved merely by MAS is not sufficient to enable accurate determination of site-specific structural information. The MQMAS method described previously must therefore be employed to improve resolution and then combined with the spectral editing power of CP. In this way, the combination of MQMAS with CP – more precisely the initial method proposed and termed CP-MQMAS [47], is used for 19 F and 27Al spin pair. The efficiency of this method was tested on a fluorinated triclinic chabazite-like AlPO4 aluminophosphate (AlPO4-CHA), (AlPO4)3(C4H10NO)F, which contains three Al sites whose NMR parameters were previously determined [48] (Table 5.1). The CP-MQMAS method involves CP from I to the quadrupolar S spin single-quantum central-transition coherences. Then these are converted (by the use of a central-transition selective 90 pulse) into a population difference across the Zeeman mS ¼ – 1/2 eigenstates from which MQ coherence is then generated for the following MQMAS experiment. This ensures that MQ coherences are generated only for spins that have a dipolar coupling to I (i.e., an editing of the spectrum occurs). It is also possible to combine REDOR with MQMAS. Two approaches, MQ-t2-REDOR [49] and MQ-t1–REDOR [50], were presented and ascertained on AlPO4-CHA. An easier experiment, termed MQMAS with dipolar dephasing (DDMQMAS), was also proposed [51] and tested on the same compound. Finally, as previously mentioned, the combination of MQMAS and TEDOR is also possible (in an analogous manner to that of MQMAS and CP) as recently shown, by Amoureux et al., on AlPO4-CJ2 [46]. 5.2.3.4

2D SLF Experiments

To monitor the effective aluminium-fluorine dipolar couplings in cryolite, as a function of temperature, amplified 2D separated local field (SLF) NMR experiments under the action of fast MAS (which is in turn desirable for the simple elimination of the homonuclear 19 19 F- F couplings) were recorded by Kotecha et al. [52]. In such SLF MAS experiments rotor-synchronized pulses were applied to achieve a net heteronuclear dipolar evolution with variable amplification factors xN of the 27Al-19F interaction along the indirect domain (x2 SLF, x4 SLF, and x8 SLF), followed by observation of aluminium’s central-transition {19F}-decoupled evolution along the direct domain [53]. 5.2.3.5

2D MAS CP-HETCOR Experiments

Heteronuclear correlation (HETCOR) solid-state NMR spectroscopy [54] has been widely used to provide information on the spatial proximity of different nuclei in complex spin systems. When more than one distinct I and/or S spin species is present, a 2D correlation experiment enables the detection of dipolar couplings between specific distinct I–S pairs. Typically, the experiment consists of a 90 pulse that creates transverse I spin magnetization, which evolves for a time t1 before it is transferred to spin S, usually via CP. The S spin FID is then detected in t2. 2D Fourier transform yields a 2D spectrum, with the appearance of cross peaks between individual I and S resonances from spins which are dipolar coupled.

Short and Medium Range Order in Al-based Fluorides

151

These 2D HETCOR experiments are also feasible when either I or S is a half-integer quadrupolar nucleus [55,56]. The spectra are able to provide information unavailable by 1D methods due to the increase in resolution obtained from the presence of a second spectral dimension. They are particularly informative if at least one of the two dimensions of the 2D spectrum is well resolved as shown on b-BaAlF5 (Figure 5.8) [57]. 19F-27Al MAS CP-HETCOR experiment was also applied in zeolite [44] and chiolite [58]. Other results will be described in Section 5.3.

(a)

(b) Al proj. Al1

19

1 cross section Al1

2

–140

3

4 5

cross section Al2

6

8

–120

7 8

F projection

9

10

9

–100

5

7

4

6

–120

–140

3 2

1

–160

19

F SQ dimension (ppm)

–100

F SQ dimension (ppm)

–160

Al2

10

10

0 27

–10 –20 –30 –40 –50 –60 Al SQ dimension (ppm)

Figure 5.8 (a) 2D 19F-27Al MAS (25 kHz) CP-HETCOR correlation spectrum of -BaAlF5. Top spectra are the full projection onto the 27Al dimension and simulation of Al1 and Al2 resonances. Dash lines indicate position where the cross-sections of Al1 and Al2 were extracted from. Right spectrum is the full projection in the 19F dimension, on which lines are labelled. (b) Selected cross-section of Al1 and Al2 and full projection in the 19F dimension of the CPHETCOR spectrum. 19F lines are labelled. Dash lines indicate 27Al-19F correlations [57]. Phys. Chem. Chem. Phys., 11, 950–957 (2009). Reproduced by permission of the PCCP Owner Societies Copyright (2009) Royal Society of Chemistry

5.2.3.6

19

F DQ-SQ MAS Correlation Experiments

Homonuclear fluorine spatial proximities can be evidenced through 2D 19F DQ-SQMAS correlation NMR experiments [59] as shown in oxyfluoride [60] and in fluoride materials [57,61–63] leading to the assignment of 19F MAS NMR spectra. In these examples, the 19 19 F- F homonuclear interaction, which is averaged out by fast MAS, is reintroduced using the back-to-back (BABA) recoupling sequence [64]. In a 2D 19F MAS DQ-SQ NMR spectrum, fluorine atomic proximities between inequivalent F sites are revealed by paired cross-correlation peaks appearing at the individual chemical shifts of the two dipolarcoupled nuclei in the SQ direct dimension and the sum of them in the indirect DQ dimension, while the proximity between two equivalent F sites is disclosed by a single auto-correlation peak located on the DQ diagonal (with a slope of 2) of the 2D spectrum

152

Functionalized Inorganic Fluorides

(Figure 5.9). Assuming that the 19F multi-spin homonuclear dipolar interaction is averaged out by the fast MAS frequency (such that the recoupled spin system can be considered as an ensemble of simple spin pairs) and using short DQ excitation/reconversion periods, the intensity of the cross-peaks in the DQ-SQ rotor-synchronized spectrum is then expected to be proportional to the number of spin pairs and to D2t2, with D the dipolar coupling constant and t the recoupling time [57 and references therein]. Advantages and limitations of this dipolar-based 2D NMR experiments for 19F MAS spectra assignments of fluoroaluminates were recently presented [57] and will be discussed in Section 5.4.1.3.

9

10

8

7

5

6

4

3 1

2

10-9 2-1

–300

3-1 4-1

5-1 7-4

6-2

8-1

5-2 4-3 6-1

3-2

7-1

–275

7-2 7-3 9-1 8-5 9-2 6-5

9-3 7-6

–250

9-4 9-5

10-2

19F

DQ dimension (ppm)

8-2 8-3 8-4

5-5

6-4

6-6

8-7

10-3 10-4 10-5 9-8 10-6 10-7

–225

10-8

–90

–100

–110 19F

–120 –130 SQ dimension (ppm)

–140

–150

Figure 5.9 2D 19F DQ-SQ MAS (30 kHz) NMR correlation spectrum of -BaAlF5. Top spectrum is the full projection in the 19F SQ dimension, on which lines are labelled. The DQ diagonals (with a slope of 2) of the 2D spectrum on which autocorrelation peaks appear are indicated by the dash lines. Paired cross-correlation peaks are indicated by horizontal solid lines [57]. Phys. Chem. Chem. Phys., 11, 950–957 (2009). Reproduced by permission of the PCCP Owner Societies Copyright (2009) Royal Society of Chemistry

Short and Medium Range Order in Al-based Fluorides

5.3

153

Application to Functionalized Al-Based Fluorides with Catalytic Properties

In this section we report on structural information obtained by high resolution solid state NMR on various aluminium fluoride phases and fluorinated aluminas and zeolites, which are extensively studied for their catalytic potential. The change in the coordination environment of aluminium on fluorination and the high electronegativity of fluorine atom are reported to be responsible for the observed modification of the acidity, and thus catalytic activity, of these materials [65–67]. Then, identification of the bulk and surface species is critical for the understanding of the catalytic properties. In the studies discussed in the following, assignment of 19F and 27Al NMR resonances of fluorinated aluminas and zeolites was achieved on the basis of model compounds such as AlF3 and AlF3 3H2O [68–69].

5.3.1

Crystalline Aluminium Fluoride Phases

AlF3 forms a rich variety of phases. All their structures are based on corner connected AlF6 octahedra. Two of them were studied by solid state NMR: -AlF3 (27Al [14,68], 19F [68,70]) and b-AlF3 (27Al and 19F [68]). Table 5.1 gathers structural information and 19F and 27Al NMR parameters. For these two polymorphs, NMR studies provide 19F iso values typical of fluorine atoms bridging two aluminium atoms (BF) and 27Al iso values typical of AlF6 octahedra. Two metastable phases of aluminium trifluoride (b- and -AlF3) and milled -AlF3 are known as suitable catalysts [71–73]. Crystalline AlF3 phases in general suffer from a low surface area (<65 m2.g1) compared to oxide catalysts [67,68,72,74].

5.3.2

19

F Isotropic Chemical Shift Scale in Octahedral Aluminium Environments with Oxygen and Fluorine in the First Coordination Sphere

Chupas et al. refined the structure of b-AlF3 3H2O. It consists of chains of corner-linked AlF4(H2O)2 octahedra held together by hydrogen-bonded interstitial water molecules. They also studied this compound by 19F and 27Al NMR [75] (Table 5.1). AlPO4-CJ2, contains isolated pairs of corner-connected AlO6xFx (x ¼ 1 and 2) octahedron and AlO4(OH,F) trigonal bipyramid [76]. Thus, 19F iso values for BF atom in an AlO5F or AlO4F2 environment and for nonbridging fluorine (NBF) atoms (i.e. fluorine atoms bound to an aluminium atom) in AlO5F and AlO4F2 environments were determined [77] (Table 5.1). Dumas et al. reported a 19F iso value for a BF atom with an AlO5F environment in a synthetic analogue of mineral minyulite K[Al2F(H2O)4(PO4)2] [78], and Simon et al. reported a 19F iso value for BF and NBF atoms in the average environment AlO3.6F2.4 in a fluoroaluminophosphate templated with 1,3 diaminopropane: [N2C3H12]Al2(PO4) (OHx, F5x) (x 2) [79] (Table 5.1). These data compiled by Chupas et al. [75] provide a partial iso scale for 19F in octahedral aluminium environments with oxygen and fluorine in the first coordination sphere, AlO6xFx.

Table 5.1 Compounds or solutions, state, fluoroaluminate species, bridging (BF) or not bridging (NBF) fluorine atom, 19F isotropic chemical shifts (ppm), 27Al isotropic chemical shift (ppm) and references for Al-based fluorides Compound

Statea

[CN]

-AlF3 -AlF3 -AlF33H2O

cr. cr. cr.

[6]

AlPO4-CHA

cr.

AlPO4-CJ2

cr.

KAl2F(H2O)4(PO4)2 [N2C3H12]Al2(PO4) (OHx, F5–x) (x2)

cr. cr.

AlF39H2O

cr.

-AlF33H2O Fluorinated -Al2O3 at 300C

cr. cr.

Fluorinated -Al2O3 at 400C

cr.

am. Fluorinated H-SSZ-32 (HS AlF3) Pristine and Zn2þ-impregnated aluminas cr. or am. Fluorinated aluminas

cr.

Al species

AlF6 AlF6 [6] AlF4(H2O)2 [6]

[6]

AlO4F2 AlO4 [6] AlO5F or [6]AlO4F2 [6] AlO5F [6] AlO4F2 [5] AlO4(OH,F) [6] AlO5F [6] AlO2OH1.6F2.4 in average [6] AlF3(H2O)3 in average

[4]

[6]

AlF3(H2O)3 in average AlF6 [6] AlFxO6–x (1<x<5) [6] AlF6 (similar to -AlF3) [6] AlF6 [6] AlF6 [6] Al(OH)xF6–x (1<x<6) [6] AlF6 [6] AlO6–xFx (1<x<6) [6]

BF or/and NBF

iso 19Fb

BF BF BF and NBF NBF

–172 –172c –145 to –160d

BF NBF NBF BF BF BF and NBF NBF

–115 –121 –124 –115 –128 –140e

NBF

–147.9 –168 –146 –173

BF BF BF BF NBF

–149.5

–171 –172 –140 to –160 –110 to –165 –180 to –220

iso 27Al

Ref.

–16 –15 –13

[14,68] [68] [75]

–4.5 43.9 and 47.6

[48] [76,77]

–8 17 –2

[78] [79]

–2.6, 3.8 and 5.8 –2.1 –17

[80] [68]

–16

[68]

–18 –16 0 to –15 –16 10

[81] [75]

[80]

[82]

Fluoroaluminate complexes in aqueous solution -Al(OH)3(s) and -AlOOH dissolved in fluoride solutions Fluorinated "-Keggin Al13O4(OH)24(H2O)127þ polycation Fluorinated -alumina

sol.

[6]

AlFx(H2O)6–x

NBF

sol.

[6]

AlFxO6–x

cr.

[6]

AlFxO6–x

BF NBF NBF

–156.9 (x ¼ 1) to –152.8 (x ¼ 5) –131 –142 –132 and –134

cr.

[6]

BF BF BF BF BF BF (3 Al) BF (2 Al) NBF BF BF NBF BF BF NBF NBF NBF NBF

–8/C6F6 33/C6F6 20/C6F6 9/C6F6 –172 –132 to –135 –145 to –149 –161 to –163 –164 0/C6F6 –40/C6F6 2/C6F6 –169 –204 –176 –188 –200

AlF6 AlO3F3 [6] AlO4F2 [6] AlO5F [6] AlF6 (in AlF3xH2O) [6] AlO5F [6] AlO5F [6] AlO5F [6] AlF6 [6] AlF6 [6]

Fluorinated -alumina

cr.

HS-AlF3 ACF, AlCl0.13F2.87

am. am.

am-AlF3 ABF, AlBr0.13F2.87

am. am.

[6]

NaF-AlF3 melts

liq.

[6]

F-bearing aluminosilicate glasses

am.

[1,8-bis (dimethylamino) naphthaleneHþ][AlF4–] [2,4,6-trimethylpyridin-Hþ][AlF4–]

cr. cr.

[4]

[N(CH3)4þ][AlF4–]

cr.

[4]

[6]

AlF6 AlF6

AlF6 AlF5 [4] AlF4 [6] AlO6F6–x [5] AlF5 [4] AlF4 or [4]AlO4–xFx [4] AlF4 [5]

[6]

AlF4 AlF6 AlF4

NBF NBF BF NBF NBF

[84] [85] [86] –17 –5 to 5 –5 to 10 5 to 10 –16 6 6 6 –14 –13 –19

[87]

[89,90] [91] [91] [92]

4 21f 38 –4 22 56 –194.2 (0.002 M in 49.2 (0.002 M CD3CN) in CD3CN) –188 –155 (broad) –190.1

[69]

48.9

[93,94] [95] [96] [96] [97,98]

(continued overleaf )

Table 5.1

(continued )

Compound

Statea

[CN]

[N(CH3)4þ]2[AlF52–] Fluorinated zeolite NaX Fluorinated zeolite HY NaF-AlF3-Al2O3 melts

cr. cr. cr. liq.

[5]

Al species

AlF5 AlO4–xFx (x1) [4] AlO4–xFx (x1) [4] AlOF3 (Al2OF6) [4] AlO2F2 (Al2O2F4) [6] HS -AlF2.6(OH)0.4 (derived from -AlF3) cr. AlF6 and [6]AlF5(OH) [6] AlF4(OH)2 [6] AlF3(OH)3 -AlF2.6(OH)0.4 treated at 873 K am. shell and [6]AlF6 and [6]AlF5(OH) [6] cr. core AlF4(OH)2 [6] AlO6 [5] AlO5 [4] AlO4 [6] HS pyrochlore aluminium cr. AlF6, [6]AlF5(OH) and [6] hydroxyfluoride: AlF4(OH)2 [6] AlF1.8(OH)1.20.3H2O AlF3(OH)3 [6] AlF2(OH)4 [6] AlF(OH)5 [5] Al [4] Al [6] Al0.82&0.18F2.46(H2O)0.54 cr. AlF6 [6] AlF5(H2O) [6] AlF4(H2O)2 [6] AlF3(H2O)3

a

[4]

crystallized (cr.), amorphous (am.), liquid (liq.), in solution (sol.). iso/CFCl3 except other indication; iso C6F6/CFCl3 ¼ 166.4 ppm. Mean value, four F sites. d At least six lines, two F sites (Al-F-Al bridges and disordered terminal equatorial positions). e Mean value, three (OH,F) sites. f An average value between 27Al iso values observed for AlF4 and AlF6 species was assumed. b c

BF or/and NBF

iso 19Fb

NBF

–161.2

23.5 47 50 50 58.5 –15.5 –11.7 –9.5 –15.6 –11.6 6.4 34.0 62.3 –12.5

–173.0 –169.4 –155.2 –152.5

–8.7 –4.2 1.8 30.9 56.1 and 64.6 –17.9 –14.4 –11.7 –9.4

–203 –173 and –182 NBF NBF BF BF BF BF BF

BF

–173.5, –167.6 and –166.7 –174.1, –173.3 and –162.3

BF BF BF BF of AlF6 BF NBF NBF

iso 27Al

Ref. [98] [99] [100] [93,94] [8] [24]

[20]

[25]

Short and Medium Range Order in Al-based Fluorides

157

More recently, - and b-AlF3 3H2O and AlF3 9H2O were studied by Kemnitz et al. [80] (Table 5.1). The structures of AlF3 9H2O and -AlF3 3H2O both consist of isolated octahedral AlF6x(H2O)x complexes with x ¼ 3 on average. All these studies show that both 19F and 27Al iso of AlO6xFx environments increase with the oxygen content. These findings are essential for the NMR investigations of disordered Al-based fluoride materials discussed in the following.

5.3.3

Fluorinated Aluminas and Zeolites, HS AlF3

In situ X-ray diffraction (XRD) and NMR methods were used to follow the structural changes that occur during the dismutation reaction of CHClF2 over -alumina [68]. 19 27 F/ Al CP experiments of -Al2O3 activated at 300 C showed that AlF3 had already begun to form at this temperature. By 400 C, resonances from a phase that resembles -AlF3 dominate both the 19F and 27Al NMR spectra of the used catalyst (Table 5.1). Considering that heterogeneous reaction rates scale with surface area, new synthetic routes to HS aluminium fluoride are currently being sought. The fluorination of H-SSZ-32 zeolite (H1.3[Al1.3Si22.7O48]) with NF3 plasma led to the formation of a HS AlF3 phase (190 m2 g1) [81] (Table 5.1). It was shown by 27Al NMR that the AlF6 octahedra in this amorphous material are greatly distorted, even in comparison with the catalytically active b-AlF3. The structures of a series of pristine and Zn2þ-impregnated aluminas following fluorination with HF were also investigated using both solid-state NMR and X-ray powder diffraction methods [75]. 19F/27Al HETCOR spectra for samples with different fluorine and aluminium environments show that the fluorine atoms with iso between 140 and 160 ppm are correlated with the aluminium sites with iso between 0 and 15 ppm. The 19F iso values become more negative as the 27Al iso shift toward more negative values. For example, 19F iso of 150 and 160 ppm correlate with 27Al iso shifts of 5 and 15 ppm, respectively. This is consistent with a later work on the mechanism for the fluorination reaction of CHClF2 on alumina determined by various 1D and 2D solid-state NMR methods [82]. Moreover an additional group of 19F resonances between 180 and 220 ppm (Table 5.1), observed on fluorinated aluminas for the first time, was assigned to NBF atoms. This range of 19F iso values for NBF atoms seems, at first sight, surprising compared to the values presented in Table 5.1 for this kind of fluorine atoms (120, 160 ppm). Furthermore, the explanation given by Chupas et al. [82], referring to the fluorine atoms bonded to a single aluminium in Na5Al3F14 showing iso values in the -200 ppm region [3], is not convincing because in this compound, these NBF atoms have several sodium atoms in their neighbouring. Nevertheless, the so-called superposition model (see Section 5.4.1.1) for the 19F isotropic chemical shift [1,16,70] shows that iso increases with the number of cations surrounding the fluorine ion (the F-M distances remaining equal). In this way, the values gathered in Table 5.1 for NBF atoms (120, 160 ppm) suggest that these NBF atoms are H-bonded to hydrogen atoms from hydroxyl or water groups. This assumption is strengthened by several studies on fluoroaluminate complexes, AlF6x(H2O)x, in aqueous solution [83,84], fluoride-promoted dissolution of bayerite b-Al(OH)3(s) and boehmite -AlOOH [85] and fluorination of the e-Keggin Al13 polycation [86] (Table 5.1). It was also shown by Chupas et al. [82] that the initial stages of fluorination lead to the formation of terminal F-Al groups and that five-coordinate aluminium sites are initially

158

Functionalized Inorganic Fluorides

consumed during the fluorination reaction. Earlier, Fischer et al. [69] and Zhang et al. [87] tentatively identified fluorine sites at the surface of fluorinated -alumina (by aqueous solutions of NH4HF2 and NH4F respectively). Obviously, despite the use of 19F/27Al HETCOR experiments [69], the assignments of the 19F NMR lines to various AlO6xFx (1 £ x £3) species [69] or to various fluorine atoms of AlO5F species (nonbridging, bridging two or three aluminium atoms) [87] disagree with above mentioned studies (Table 5.1). Finally, the thermal degradation of a perfluoropolyether on the surfaces of -alumina and kaolinite was studied by Denkenberger et al. [88]. If the formation of aluminium oxy-fluoride species is obvious, the identification of these species is questionable since it is based on the work of Fischer et al. [69]. In 2003, Kemnitz and coworkers reported a new nonaqueous synthesis route to HS amorphous AlF3 [89]. 27Al and 19F [90] NMR spectra indicate that this compound contains AlF6 octahedra (Table 5.1).

5.3.4

Aluminium Chlorofluoride and Bromofluoride

HS-AlF3 is an extremely strong Lewis acid, comparable with aluminium chlorofluoride (ACF: AlClxF3x, with x ¼ 0.05–0.30). Structural insights on ACF and an amorphous phase of AlF3 (am-AlF3), which is known to emerge during heating of -AlF3 3H2O to 200 C under a self-generating atmosphere, were obtained by 19F MAS and 27Al SATRAS NMR [91]. It was shown that a higher degree of disorder exists in ACF than in am-AlF3 and that small amount (4–5%) of NBF atoms and nearly regular AlF6 octahedra are present in the ACF sample (Table 5.1). The synthesis and the characterization of amorphous aluminium bromide fluoride (ABF: AlBrxF3x, with x ¼ 0.13), which is very similar to ACF, were also reported [92] (Table 5.1).

5.3.5

Pentahedral and Tetrahedral Aluminium Fluoride Species

Whereas all fluoroaluminate complexes, AlF6x(H2O)x, observed in aqueous solution, for all pHs less than 8, and for all [F]/[Al] ratios, are hexacoordinated, with an octahedral geometry [83], penta and tetracoordinated fluoroaluminate complexes undoubtedly occur in the liquid NaF-AlF3 melts, as shown by high-temperature 27Al and 19F NMR spectroscopies [93,94] (Table 5.1), and were observed in the solid state. In 1991, Kohn et al. gave evidence for four, five- and six-coordinated aluminium fluoride complexes in F-bearing aluminosilicate glasses (Table 5.1). The [5]Al was thought to be present in AlF52 complexes, a previously unknown species [95]. In 1993, Herron et al. showed that the AlF4 anion can be solubilized in organic solvents and crystallized as organocation salts and is tetrahedral [96,97] (Table 5.1) and in 2003, Groß et al. synthesized and characterized the AlF52 anion as [N(CH3)4þ]2[AlF52] [98] (Table 5.1). Tetrahedral aluminium oxyfluoride (AlO4xFx) species were produced by the dehydrofluorination reaction of hydrofluorocarbon-134 over basic faujasite zeolites [99] and by the fluorination of zeolite HY by NH4F [100] (Table 5.1). A multinuclear NMR study of high-temperature NaF-AlF3-Al2O3 melts suggests the existence of at least two different tetrahedral oxy-fluoride species Al2OF62 (AlOF3) and Al2O2F42 (AlO2F2) in the liquid in addition to the widespread AlF4, AlF52, and AlF63 species [94]. Finally, pentahedral oxyfluoride species were

Short and Medium Range Order in Al-based Fluorides

159

evidenced in hign peralkaline aluminosilicate glasses [101] and in the Na2B4O7-Na3A1F6TiO2 system [102] as well as pentahedral fluoride or oxyfluoride and tetrahedral oxyfluoride species in aluminosilicate and fluoride-containing glasses [103].

5.3.6

Nanostructured Aluminium Hydroxyfluorides and Aluminium Fluoride Hydrate with Cationic Vacancies

Recently, Dambourbet et al. developed a new route enabling the synthesis of Al-based fluoride materials presenting nanosized crystallized (HS) particles (see Chapter 2 of this book). Microwave-assisted synthesis was applied and nanostructured aluminium hydroxyfluorides derived from b-AlF3 [8], nanostructured Al-based fluoride-oxide materials with a core-shell morphology [24], HS pyrochlore aluminium hydroxyfluoride [20] and an aluminium fluoride hydrate with cationic vacancies [25] were prepared and structurally characterized. 27Al high-field NMR spectroscopy was used to probe and quantify the various AlF6x(OH)x, AlOx or AlF6x(H2O)x environments occurring in these compounds (Figure 5.10). In these four studies, the assignments of the 27Al NMR contributions

Al0.82

0.18F2.46(H2O)0.54

Exp. Cal. C. T. S. T.

AlF1.8(OH)1.2, 0.3H2O

Exp. Cal. C. T. S. T. 20

10

0 27Al

–10 –20 frequency (ppm)

–30

–40

Figure 5.10 Experimental and calculated central transitions of the 27Al MAS (30 kHz) NMR spectra of Al0.82&0.18F2.46(H2O)0.54 and pyrochlore AlF1.8(OH)1.2 0.3H2O recorded at 17.6 T. The fitting, achieved using four contributions, takes into account the N ¼ 0 band of both the satellite transitions (S.T.) <3/2> and the central transition (C.T.) <1/2>

160

Functionalized Inorganic Fluorides

(Table 5.1) were achieved assuming that the higher the number of oxygen atoms in AlF6x(OH)x or AlF6x(H2O)x species, the higher the 27Al chemical shift. For b-AlF3x(OH)x [8], NMR investigations supported the localization and the statistic distribution of the OH groups in F1 and F2 sites and allowed stating more precisely the chemical compositions. In the Al-based fluoride-oxide materials exhibiting a core-shell morphology, 19F and high field 27Al NMR contributed to identify the core (crystallized) part of the material as closely related to b-AlF3. The amorphous shell contains some AlO6, AlO5, and AlO4 species that were identified and quantified [24]. The framework of pyrochlore aluminium hydroxyfluoride is built of AlF6x(OH)x species. A random distribution of F atoms and OH groups was shown to occur on the 48f sites [20]. Finally, in the compound containing aluminium vacancy Al0.82&0.18F2.46(H2O)0.54, unprecedented in the Al-based fluoride chemistry, the species AlF6 and AlF6x(H2O)x with x ¼ 1, 2, 3 were identified and quantified. The quantification led to the formula Al0.82F2.22(H2O)0.24&0.18F0.24(H2O)0.30, which gives independent insights of both the Al3þ ions and vacancy environments and shows the occurrence of unusual H2O molecules surrounded by two cationic vacancies. Actually, the vacancies are mainly surrounded by water molecules but also by a low content of fluoride ions as also evidenced by 19F NMR showing the occurrence of NBF atoms [26]. 5.3.7

d iso Scale for 27Al and 19F in Octahedral Aluminium Environments with Hydroxyl and Fluorine in the First Coordination Sphere

Following the studies mentioned above, pyrochlore AlFx(OH)3x H2O with variable Fcontent was investigated by 19F MAS NMR and 27Al NMR at different magnetic fields up to 21.1 T, by Ko¨nig et al. [104,105]. In AlF6x(OH)x species 27Al iso increases nearly linearly with x (the studied samples are too numerous to be reported in Table 5.1) [105]. The 19F MAS NMR experiments show that 19F iso increases also nearly linearly with x for the same species [102]. Dehydration of the samples revealed the influence of the water molecules on 19F iso [104] and allows determining trend analysis for 19F in proton-poor substances. To conclude this section, we would like to indicate that [4]Al, [5]Al and [6]Al coordination are easily discriminated due to their different 27Al chemical shifts. Moreover for crystallized disordered compounds with homogeneous composition and known structure, high field NMR experiments allow differentiating species with the same coordination number. On the contrary, for fluorinated aluminas and zeolithes and amorphous compounds, despite proximities between 19F and 27Al clearly revealed by correlation experiments, the differentiation of such species remains difficult.

5.4

Alkali and Alkaline-earth Fluoroaluminates: Model Compounds for Modelling of NMR Parameters

This section is devoted to the study of alkali and alkaline-earth fluoroaluminates. These systems were selected because they include numerous crystalline phases with known

Short and Medium Range Order in Al-based Fluorides

161

structures. First, the 19F NMR lines are assigned to their crystallographic sites using semiempirical or ab initio numerical methods that provide 19F iso parameters with good accuracy. We also show that 19F DQ-SQ MAS NMR correlation experiments (Section 5.2.3.6) may be a complementary tool for these assignments. The second part deals with the assignment of the 27Al NMR lines, through the modelling of the 27Al NMR quadrupolar parameters, using ab initio methods. The calculations allow the structure and electronic characterization of the studied compounds.

5.4.1

19

F NMR Line Assignments

Numerous studies have suggested a relationship between 19F iso and the local environment of the fluorine atoms. The approach that is often used is based on the intuitive idea that similar chemical shift values indicate similar structural environments. By comparison with the chemical shift values obtained in basic well-known fluorides having a single crystallographic site, NMR lines of a crystalline compound can then be assigned to different fluorine environments through their position on the spectrum [106–113]. When fluorine sites have different multiplicities, the relative intensities of the resonances also give constraints for their assignment. Nonetheless, complete unambiguous assignment of complex 19F solid-state MAS NMR spectra often remains challenging. One solution is to calculate the iso from structural data, either with semi-empirical models or with ab initio codes. 5.4.1.1

Superposition Model for 19F iso Calculations

Bureau et al. [70] proposed a semi-empirical superposition model: the 19F iso is considered as a sum of one constant diamagnetic term and several paramagnetic contributions from the neighbouring M cations depending on the M-F distance, d. The calculation of iso is performed using the following formula: iso = C6 F6 ¼ 127:1 1 1 with l ¼ l ; 0 exp ½l ðd d0 Þ d0 is taken as equal to the F-M distance in the related basic fluoride (MFn) l,0 determines the order of magnitude of the cationic paramagnetic contribution to the shielding and was deduced from measurements in the related basic fluoride. l describes the behaviour of the paramagnetic contribution with the F-M distance. This model was satisfactorily tested on fluorides of increasing complexity [70]. In order to improve this model, we recorded and reconstructed high-speed MAS 19F NMR spectra for 13 compounds from the BaF2-AlF3, CaF2-AlF3, BaF2-CaF2-AlF3 [1,2] and NaF-CaF2-AlF3 systems [16] which led to the determination of 95 iso values for 19F in various environments. A first attribution of the NMR lines was performed for eight compounds from the BaF2AlF3 and CaF2-AlF3 systems using the superposition model as initially proposed by Bureau et al. [70]. The phenomenological parameters of this model (d0, l,0 and l) were then refined to obtain a better agreement between experimental and calculated values, i.e. to improve the NMR line assignment. Satisfactory reliability was reached with a rootmean-square (RMS) deviation between calculated and measured values equal to 6 ppm (Figure 5.11). The refined parameters were successfully tested on -BaCaAlF7. Finally,

162

Functionalized Inorganic Fluorides

the iso ranges were defined for the BF, NBF, and ‘free’ (not embedded into any AlF6 octahedron) fluorine atoms encountered in the investigated binary systems. In compounds whose structures are unknown, the fluorine surroundings can then be deduced from the NMR line positions. For example, in b-CaAlF5, this confirmed the proposition that this compound is isostructural with CaFeF5. Since then, this assumption was ascertained by the refinement of the structure [2]. Such an approach was also applied to BaF2-CaF2-AlF3 glasses [114]. Similar considerations were applied to relate the 18 iso values measured for three compounds from the NaF-CaF2-AlF3 system to their crystallographic sites [16]. However, uncertainties on calculation results sometimes are larger than the experimentally measured 19F iso difference between two distinct resonances [1,16], questioning the reliability of some assignments.

250

calculated 19F δiso (ppm)

200

150 AlF3

100

BaF2 CaF2 α-CaAlF5 Ca2AlF7

50

α-BaAlF5 β-BaAlF5 γ-BaAlF5

0

Ba3Al2F12 β-Ba3AlF2 Ba3AlF9-Ib

–50 –50

0

50 100 150 experimental 19Fδiso (ppm)

200

250

Figure 5.11 Isotropic chemical shift values calculated with refined phenomenological parameters versus experimental ones for compounds from BaF2-AlF3 and CaF2-AlF3 binary systems. The linear regression (solid line) and the iso, calc ¼ iso, exp curves (dashed line) are nearly superimposed [1]. Reprinted with permission from Inorg. Chem., 43, 2474–2485 (2004). Copyright 2004 American Chemical Society

5.4.1.2

Cluster Based ab initio Calculations of 19F iso

Ab initio calculations of 19F iso have also been carried out, using gauge-independent atomic orbitals (GIAO) at the DFT level, on -AlF3 [115–117] chiolite [118], various [4] Al, [5]Al and [6]Al oxyfluorinated species in order to study possible Al-F

Short and Medium Range Order in Al-based Fluorides

163

bonding environment in fluorine-bearing sodium aluminosilicate glasses [118,119] and on aluminium fluoro-complexes in order to characterize extraframework aluminium in H-mordenite dealuminated with (NH4)2SiF6 [120]. Clusters centred on the studied fluorine atoms mimic the crystalline structures. Liu et al. [118,119] and Cai et al. [115,116] clearly showed that the calculated isotropic shielding iso values not only depend on the basis sets employed but also on the number of atoms included in the calculation. Body et al. [121] applied such ab initio calculations to a large number of inorganic crystallized fluorides including AlF3, alkali and alkaline-earth fluoroaluminates, with the goal of working out a tractable method, in the choice of the basis sets and in the generation of the clusters. In total, chemical shifts were calculated for 122 fluorine sites. In these compounds, for the clusters without barium for which the quality of the basis sets is poor, the ab initio method led to a RMS equal to 22 ppm, which is quite a nice result keeping in mind that the 19F iso range is larger than 200 ppm. However, such RMS value does not allow performing reliable assignments, and the cluster build up remains empirical for some fluorine atoms. One way to overcome the limitation of the cluster method is to take into account the crystal periodicity in the chemical shift computation through the use of software that includes the periodic boundary conditions [122].

5.4.1.3

19

F DQ-SQ MAS Correlation Experiments

As shown by Martineau et al. [57], another strategy to assign 19F NMR lines in fluoroaluminates is to perform 2D MAS NMR correlation experiments, 19F-27Al CP-HETCOR and 19F DQ-SQ. For the studied barium and calcium fluoroaluminates, only 19F-19F ˚ (corresponding to 19F dipolar couplings larger than proximities shorter than 3.6 A 2.4 kHz) were experimentally revealed by correlation peaks of significant intensities in the 2D 19F DQ-SQ MAS NMR correlation spectra. In b-BaAlF5, combination of 2D MAS 19 27 F- Al CP-HETCOR (Figure 5.8) and 19F-19F DQ-SQ NMR correlation (Figure 5.9) experiments allow complete unambiguous assignment of the ten poorly resolved resonances of same relative intensities of the 19F MAS NMR spectrum. The gain in resolution of the 19F MAS 2D spectrum compared to a 1D spectrum is evidenced, allowing distinction between the two BF resonances of Ba3Al2F12. In this compound, six of the eight fluorine resonances are assigned. Distinction between the two ‘free’ fluorine ions cannot be made since the dipolar-recoupling between their resonances and the BF ions is precluded due to the 19F iso range being too large. In -CaAlF5, the two NBF resonances cannot be distinguished because they show similar auto- and cross-correlation peak intensities despite their significantly different F-F inter-atomic distances, illustrating the strongly coupled multi-spin character of the 19F rigid spin network. Overcoming this limitation would require the use of a higher magnetic field combined with the recently accessible ultrafast MAS, which opens a new route to study numerous fluorides with strongly coupled multispin systems. For b-BaAlF5 and Ba3Al2F12, line assignments obtained from 19F DQ-SQ MAS correlation experiments [57] question those established from these 19F iso calculations [1]. This demonstrates that such experiments are essential for correct line assignments and will be also useful to assess 19F iso calculation results.

164

5.4.2

Functionalized Inorganic Fluorides 27

Al Site assignments, Structural and Electronic Characterizations

The electric field gradient (EFG) is a ground state property of solids that sensitively depends on the asymmetry of the electronic charge density near the probe nucleus. The EFG is defined as the second derivative of the electrostatic potential at the nucleus position written as a traceless tensor. A nucleus with a nuclear spin number I 1 has a nuclear quadrupole moment (Q) that interacts with the EFG which originates from the nonspherical charge distribution surrounding this nucleus. This interaction 3eQVzz and the asymmetry determines the nuclear quadrupolar frequency Q ¼ 2I ð2I 1Þh V V parameter Q ¼ yyVzz xx . The Vii are the eigenvalues of the EFG tensor with the convention jVzz j jVxx j Vyy , e is the electron charge, I is the nuclear spin quantum number, and h is the Planck constant. 27 Al SATRAS and MQ-MAS spectra were recorded and reconstructed for the compounds presented in Section 5.4.1.1 and for -AlF3, Na3AlF6 and Na5Al3F14 leading to the precise determination of the 27Al NMR parameters (iso, Q, Q) [2,14–17]. The main finding of the study of the crystalline compounds from BaF2-AlF3, CaF2-AlF3 and BaF2-CaF2-AlF3 systems, is the dependence of the Q on the type of AlF63 octahedron connectivity [17]. The experimental Q values range between 75 kHz and 510 kHz for structures built up from isolated octahedra, between 560 kHz and 1250 kHz for structures built up from isolated chains of cis-connected octahedra, and between 1530 kHz and 1580 kHz for structures built up from isolated chains of trans-connected octahedra. Then, the Q value, which is related to the largest component of the EFG, could account for the way in which the AlF63 octahedra are connected, i.e. in a trans, cis or isolated manner. In order to gain a deeper understanding of the dependence of the Q values on the AlF63 octahedron network connectivity, ab initio EFG calculations of 27Al in crystalline phases were achieved. Blaha et al. [123] showed that EFG can be calculated for large infinite solids described within periodic boundary conditions, using the full-potential linearized augmented plane-wave (FPLAPW) method. We achieved such DFT-based calculations of 27Al quadrupolar parameters of 23 aluminium sites in 16 fluoroaluminates [124] using the WIEN2k package [125]. For -AlF3, -Na3AlF6, and Na5Al3F14, 27Al EFG calculations were previously performed without optimizing the structures [14]. In a study on aluminofluoride minerals, Zhou et al. determined and calculated 27Al quadrupolar parameters in the latter two compounds [126]. Their experimental parameters are very close to those published in [14]. For -CaAlF5, b-Ba3AlF9, and compounds from the NaF-CaF2-AlF3 system, the structure optimizations and 27Al EFG calculations were published in [2,15,16]. We demonstrated that accurate NMR quadrupolar parameters represent valuable experimental data for evaluation and refinement of inorganic structures when combined with high-level DFT calculations and structure optimizations by minimization of the forces acting on the nuclei [2,15,16,124]. Before optimization, the agreement between calculated and experimental values for both Q and Q was not satisfactory, as shown in Figure 5.12 for Q, the discrepancy being larger for the structures determined from powder diffraction data, usually less accurate.

Short and Medium Range Order in Al-based Fluorides

165

A really improved agreement was obtained using the optimized structures. The shifts of the atomic positions lead to reduced octahedron distortions. The optimized structures provided a reliable assignment of the 27Al quadrupolar parameters to the aluminium sites in the four studied compounds, which contain two sites with the same multiplicity [15,16,124]. The correlation between experimental and calculated EFG tensor elements (Figure 5.12) allowed the determination of the value of the 27Al nuclear quadrupole moment with an improved accuracy [124]. Moreover the DFT calculations provided the orientation of the 27 Al EFG tensors in the crystal frame and allowed a quantitative interpretation of the tensor element orientations and magnitudes in terms of electron densities and octahedron distortions. Electron density maps (Figure 5.13) suggested that the magnitude and orientation of the 27Al EFG tensors in fluoroaluminates mainly result from the asymmetric distribution of the Al 3p orbital valence electrons. For most of the aluminium sites with values of Vzz higher than 0.75 1021 V.m2, the Vzz direction were shown to be oriented along, or nearly along, two opposite Al-F bonds and the sorted Vii tensor elements are directly proportional to the three sorted dF Al F hdi radial distortions (Figure 5.14). Charge concentrations or depletions are located along the Vzz direction corresponding to the highest jdF Al F hdij (absolute value) (Figure 5.13). The

1.5e + 6

Experimental νQ (Hz)

1.0e + 6

5.0e + 5

0.0

–5.0e + 5

–1.0e + 6

–1.5e + 6 –2.0e + 21 –1.0e + 21 0.0 1.0e + 21 2.0e + 21 3.0e + 21 Calculated Vzz (V·m–2)

Figure 5.12 27Al quadrupolar frequency versus calculated Vzz (&) before and ( ) after optimization. The solid line corresponds to a linear regression Q ¼ 5.86 10–16 Vzz (R2 ¼ 0.999) after optimization (except for -AlF3 and -Na3AlF6). Error bars indicate uncertainties in quadrupolar frequency [124]. Reprinted with permission from J. Phys. Chem. A, 111, 11873–11884 (2007), copyright 2007 American Chemical Society

•

166

Functionalized Inorganic Fluorides a)

b)

F1 F3

Al F3

Ca

F1

Figure 5.13 (a) Orientation of the calculated 27Al EFG tensor in -CaAlF5. The vector lengths ˚ ) and F labels are are proportional to the magnitude of the contributions. Al-F distances (A indicated. (b) Difference electron density (D ) in the F1-Al-F3 plane in -CaAlF5. Atom labels ˚ 3. Solid, dotted, and dashed lines are indicated. The contour intervals are in units of 0.05 e.A correspond to positive, zero, and negative D , respectively [124]. Reprinted with permission from J. Phys. Chem. A, 111, 11873–11884 (2007), copyright 2007 American Chemical Society

largest EFGs are observed for isolated chains of corner sharing AlF63 octahedra involving two BF atoms. Actually, the chiolite case demonstrated that the charge density distributions around aluminium atoms mostly arise from the different nature (BF and NBF) of the fluorine atoms in the AlF63 octahedron and are not systematically related to radial distortions.

Short and Medium Range Order in Al-based Fluorides

167

3 α -CaAlF β -CaAlF

Vzz and Vxx (1021 V.m–2)

2

α -BaAlF β -BaAlF

1

γ -BaAlF β -Ba AlF Al3 α -NaCaAlF

0

Linear regr. Na Al F Ba Al F

–1

–2

–3 –0.15

–0.10

–0.05 0.00 0.05 dF–Al–F – (Å)

0.10

0.15

Figure 5.14 Vzz and Vxx EFG tensor elements vs radial distortion for Al sites with an absolute value of Vzz larger than 0.75 1021 V.m–2. The solid line corresponds to the linear regression (a ¼ 23.6 1031 V.m–3, b ¼ –0.029 1021 V.m–2, R2 ¼ 0.984) for which Na5Al3F14 and Ba3Al2F12 are not taken into account (the corresponding symbols are surrounded by circles) [124]. Reprinted with permission from J. Phys. Chem. A, 111, 11873–11884 (2007), copyright 2007 American Chemical Society

For sites with values of Vzz lower than 0.60 1021 V.m2, the AlF63 octahedra are isolated (except -AlF3) and all the Vii are oriented along the Al-F bonds. Then, the angular distortions have to be taken into account. When the radial distortions are very small, the definition of relevant indices allows correlation between angular distortions, which are the predominant factor, and magnitude and sign of the tensor elements. To conclude, this study [124] shows the strong potential of combining EFG DFT calculations and accurate quadrupolar NMR parameter measurements for the structural and electronic characterizations of crystalline inorganic materials.

5.5

Conclusion

This chapter has shown, using Al-based fluorides as examples, that current and emerging solid state NMR methods may provide valuable answers to structural questions of relevance in material science. New improvements may be expected for dipolar coupled nuclei from the combination of very fast MAS with multiple quantum and multi-dimensional experiment at very high magnetic field. The value of a coupled experimental and theoretical approach of the NMR parameters is illustrated on 19F and 27Al. It demonstrates the advantages of the use of DFT calculations to correlate these parameters to structural features. The current increase of computational power is likely to make the use of these ab initio calculations more popular to assist the study of complex disordered or amorphous material.

168

Functionalized Inorganic Fluorides

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Short and Medium Range Order in Al-based Fluorides

[21] [22] [23] [24] [25]

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6 Predictive Modelling of Aluminium Fluoride Surfaces Christine L. Bailey1, Sanghamitra Mukhopadhyay2, Adrian Wander1, Barry Searle1 and Nicholas Harrison1,2 1

Computational Science and Engineering Department, STFC Daresbury Laboratory, Daresbury, Warrington, Cheshire WA4 4AD, United Kingdom 2 Department of Chemistry, Imperial College London, Exhibition Road, London SW7 2AZ, United Kingdom

6.1

Introduction

The search for new materials with specific properties involves the synthesis and characterization of a large selection of related materials. Understanding the composition, structure and properties of any newly synthesized material is essential in this development process. Of even greater importance, however, is the understanding of how the composition of a material, its structure and its properties are all related to each other. If these relationships can be understood then predictions can be made about the modifications required to synthesize a material with desired properties. This detailed understanding of these relationships allows us to develop a conceptual framework for a given group of related materials. In general, the relationships are very hard to establish using only experimental techniques. This is because it is not possible to have complete control over the synthesis procedure or to have detailed knowledge of the structure and composition of a material, nor is it always possible to measure the key properties of a material. Theoretical modelling techniques are essential in the development of a conceptual framework; the exact

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Ó 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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composition and structure of the material investigated can be controlled and, in many cases, reliable calculations of relevant properties can be performed. In this way a detailed understanding of the relationship between structure, composition and properties of a material can be obtained. Only in recent years has the accuracy, reliability and the efficiency of ab initio calculations allowed us to study complex materials for which detailed experimental characterization is missing. In the real world, many important chemical reactions take place under high temperature and high pressure conditions. Ab initio calculations have traditionally been thought of as zero-temperature, zero-pressure techniques, however they can be combined with atomistic thermodynamics to include the effects of finite temperatures and pressures. This allows predictive modelling of surface composition and structure under realistic conditions. We apply ab initio atomistic thermodynamic modelling techniques to the study of aluminium fluorides. We predict the atomistic structure and composition of crystalline - and -AlF3 under realistic conditions. We then extend this work to make predictions of the true nature of the synthesized HS-AlF3 materials. An understanding of how the two differ may lead to improved synthesis methods that allow the production of materials with preferred characteristics. The tools and techniques used and described in this study are transferable to the study of many other materials. It is hoped that this chapter will demonstrate how recent advances in theoretical methods can be used to interpret experimental results and to predict and understand the functionality of new materials. This chapter is organized as follows. In the next section we discuss the theoretical methods used in this chapter. The geometric structure and energetics of the surfaces of - and -AlF3 are then discussed. The surfaces are characterized by investigating their ability to bind the Lewis base NH3. Hydroxylation of the surfaces and the adsorption of H2O and HF is then considered and their structure is predicted as a function of H2O and HF chemical potential. The reactivity of these partially hydroxylated surfaces is investigated by studying their interaction with CO. Finally we consider the mechanism by which the dismutation of CCl2F2 is catalysed at the surface of -AlF3.

6.2 6.2.1

Methodology Density Functional Theory

Ab initio calculations discussed in this chapter were performed using the CRYSTAL code [1]. Electronic exchange and correlation were treated using the hybrid-exchange (B3LYP) approximation [2] to density functional theory. This functional has been shown to provide reliable geometric and electronic structures and energetics in a wide range of materials [3]. Triple valence local Gaussian basis sets were used throughout. Details of these basis sets can be found elsewhere [4, 5, 6]. Structures were fully optimized, within symmetry constraints, using a BroydenFletcher-Goldfarb-Shanno (BFGS) [7] algorithm. Optimization was considered to be complete when the residual forces were below 1.0 104 Hartrees Bohr1. The

Predictive Modelling of Aluminium Fluoride Surfaces

177

binding energies of molecules were corrected, where possible, for basis set superposition error (BSSE) using the counterpoise scheme [8]. The vibrational frequencies of the molecules were computed by construction of the force constant matrix via finite differencing of the analytic gradients followed by diagonalization of the resultant dynamical matrix [9]. Transition states were calculated using the nudged elastic band algorithm [10,11].

6.2.2

Surface Free Energies

The relative stability of surfaces of different composition in contact with reactive gases is determined by the free energy of surface formation which can be calculated from first principles. The formalism in the context of metal oxides is now well established [12–15] and is applied here to the specific reaction conditions of AlF3 [16]. Modelling the AlF3 surface as a periodic slab of material, the surface free energy is given by [17] 1 1 Gslab ðT; PF2 Þ NAl mAl ðT; PF2 Þ NF mF2 ðT; PF2 Þ ðT; PF2 Þ ¼ 2A 2

(6:1)

where A is the area of the unit cell (the factor of two accounts for both sides of the slab). Gslab is the Gibbs free energy per unit cell of the slab. NAl and NF are respectively the total number of Al and F ions within the system, PF2 is the partial pressure of the gaseous F2 and mAl and mF are the chemical potentials of Al and F respectively. The condition that the bulk aluminium fluoride is in equilibrium with pure Al and gaseous F2 is

3 Gbulk ðT; PF2 Þ ¼ mAl ðT; PF2 Þ mF2 ðT; PF2 Þ 2

(6:2)

where Gbulk is the Gibbs free energy per formula unit of bulk AlF3. Equation (6.2) can be used to eliminate mAl from equation (6.1) to obtain 1 1 ðT; PF2 Þ ¼ Gslab ðT; PF2 Þ NAl Gbulk ðT; PF2 Þ ðNF 3NAl ÞmF2 ðT; PF2 Þ 2A 2

(6:3)

This methodology can be extended to model an AlF3 surface with adsorbed hydroxyl groups, H2O and HF exposed to an atmosphere containing gaseous H2O, H2 and HF. The surface free energy is then; 2

3 1 m ðT; P Þ N m ðT; P Þ N ðT; P Þ G tot Al Al tot F F2 7 1 6 slab 2 F2 ðT; PF2 ; PO2 ; PH2 Þ¼ 4 5 1 1 2A N m ðT; P Þ N m ðT; P Þ O O2 H H2 O2 H2 2 2

(6:4)

178

Functionalized Inorganic Fluorides

where Ptot is the sum of the partial pressures of the individual gases. HF and H2O are in equilibrium with their constituent atoms, hence 1 1 m ðT; PH2 Þ þ mF2 ðT; PF2 Þ¼mHF ðT; PHF Þ (6:5) 2 H2 2

1 mH2 ðT; PH2 Þ þ mO2 ðT; PO2 Þ ¼mH2 O ðT; PH2 O Þ 2

(6:6)

Equations (6.2), (6.5) and (6.6) can now be used to eliminate mAl, mF2 and mO2 from Equation (6.4) to obtain 2 ðT; PHF ; PH2 O ; PH2 Þ ¼

Gslab ðT; Ptot Þ NAl Gbulk ðT; Ptot Þ

3

7 1 6 6 ðNF 3NAl ÞmHF ðT; PHF Þ NO mH2 O ðT; PH2 O Þ 7 6 7 5 2A 4 1 ð3NAl NF 2NO þ NH ÞmH2 ðT; PH2 Þ 2

(6:7)

Treating the gaseous species as ideal gases, their chemical potentials dependence on P and T is PX o (6:8) mX ðT; PX Þ ¼mX ðT; PX Þ þ kT ln PoX This chemical potential can be referred to the athermal limit and the DFT calculations by rewriting Equation (6.8) as m0X ðT; PX Þ ¼ mX ðT; PX Þ mX ð0; PoX Þ þ EDFT ðT¼ 0Þ

(6:9)

The term in square brackets in Equation 6.9 can be obtained from thermodynamical reference tables [18], as described previously [16]. In the current study the Gibbs free energies of the slab and bulk crystal are computed at the athermal limit and their temperature dependence is ignored as it is negligible compared to that of the gaseous species. The small PV term due to the change in volume of the bulk phases is also neglected. 6.2.3

Molecular Adsorption

It is possible to extend the formulism discussed in Section 6.2.2 to calculate the extent to which a gas will chemisorb to a given surface. In equilibrium, the rate of desorption of molecules from the surface will be equal to their rate of adsorption. The ratio of the rate of adsorption, rads, to the rate of desorption, rdes, is approximately rads Dm Eads (6:10) ¼ exp rdes kT

Predictive Modelling of Aluminium Fluoride Surfaces

179

where Dm is the change in the Gibbs free energy per molecule between the gas phase and the adsorbed phase and Eads is the binding energy of the molecule to the surface. This approximation neglects any change in the Gibbs free energy of the surface after adsorption, which has previously been shown to be very small [14]. The chemical potential of an ideal gas is given by 2 3 3 2 kT 2pm kT i 5 mgas ¼ kT 4 h2 pi

(6:11)

where mi is the mass of a molecule of the gas, h is Planck’s constant and pi is the partial pressure of the gas. The chemical potential of a gas molecule (such as CCl2F2 which is discussed in section 6.7.3) also contains terms due to the internal degrees of freedom of the molecule. It can be assumed that these terms will be dominated by the vibrational degrees of freedom, which will only change by a small amount after adsorption. Dm in Equation (6.10) can, therefore, be written in terms of the chemical potential of an ideal gas. Using this approximation Dm ¼ –0.46 eV at 300 K and –1.01 eV at 600 K for CCl2F2.

6.2.4

Kinetic Monte Carlo Simulations

Kinetic Monte Carlo simulations can be parameterized from DFT calculations and used to model surface processes. We shall demonstrate how they can be used to predict the rate at which NH3 molecules desorb from a surface during a temperature programmed desorption experiment, assuming there is no readsorption of molecules to the surface. Initially a grid, containing A sites, is set up to represent the surface at time zero. When modelling the desorption of NH3 from a surface, each site on the grid represents an adsorbed NH3 molecule. The rate constant for event j occurring at site i is given by jDEði; jÞj ri ; j ¼ o exp kT

(6:12)

where 0 is the attempt frequency, k is the Boltzmann constant, T is the temperature of the system and DE(i, j) is the energy barrier associated with event j at site i, its value is also dependent on the occupancy of neighbouring grid sites. The rate constants are calculated for each grid point. The cumulative function Rm ¼

m X

rn

(6:13)

n¼1

is calculated for m ¼ 1, N where N is the total number of possible events. A random number, u [1, N], is generated and the event m that satisfies Rm1 < m < Rm is selected. The grid is updated to represent the occurrence of this event and the time is updated by u t ¼ t ln Dt (6:14) N

180

Functionalized Inorganic Fluorides

where Dt ¼ 1/RN. Dt is multiplied by ln(u/N) as this captures the stochastic nature of the time step. Updating the temperature at each time step allows heating or cooling to be modelled. Rate constants, rij, that change as a consequence of event m occurring are updated (if the temperature changes then all the reaction rates must be updated). The cumulative function is recalculated, a new random number generated and the process is repeated.

6.3 6.3.1

Geometric Structure of a and b-AlF3 Bulk Phases

The various crystalline forms of AlF3 consist of arrangements of corner sharing AlF3 octahedra. The most stable phase is -AlF3, which has a close packed corundum-like structure, as shown in Figure 6.1a. The metastable -phase has a more open structure of the hexagonal tungsten bronze (HTB) type and is shown in Figure 6.1b. The computed equilibrium lattice constant [19] for the and unit cells are slightly larger, by not more than 2%, than those observed [20, 21]. This is as expected as B3LYP calculations typically overestimate the bond lengths of ionic materials.

(a)

(b)

Figure 6.1 (a) Bulk a-AlF3. (b) Bulk b-AlF3

6.3.2

Surfaces

When a bulk crystal is cleaved to form a surface the coordination of the surface atoms is reduced and the atoms close to the surface relax to minimize the surface energy of the crystal. As a result of this, the distances between the first few layers of the crystal will change with respect to their bulk values; this process is known as surface relaxation. A more severe change of structure is the phenomenon of surface reconstruction. This involves larger (still on the atomic scale) displacements of atoms compared to surface relaxation and the periodicity parallel to the surface may change with respect to that of the bulk.

Predictive Modelling of Aluminium Fluoride Surfaces

181

F2 partial pressure at 300K (atm) 10–30

Surface Energy (Jm–2)

8 7 6

10–20

10–10

100

0F 1F 2F 3F 4F 5F 6F

5 4 3 2 1 0

–2.5

–2

–1.5 –1 F2 chemical potential (eV)

–0.5

0

Figure 6.2 The surface energies of the various terminations of the (0001) (1 2) a-AlF3 surfaces as a function of fluorine chemical potential and fluorine partial pressure at 300 K. The surfaces are labelled by the number of F ions in their surface-most layers. The 3F surface is shown in Figure 6.3

Al(4)

Al(5)

Figure 6.3 The structure of the stoichiometric (0001) (1 2) a-AlF3 3F surface. The F ions are represented by large spheres and the Al ions by small spheres

Several different terminations of the -AlF3(0001) (1 2) surface have been relaxed and their surface energies are plotted in Figure 6.2, as a function of fluorine chemical potential and partial pressure at 300 K (using the methodology described in Section 6.2.2). The surface in which the top most plane consists of three F ions is the most stable at all fluorine chemical potentials. This structure is shown in Figure 6.3. At zero chemical potential of F2 (i.e. low temperature and very high F2 partial pressure) the surface energy of the 5F surface is identical to that of the 3F surface. This is because the optimized 5F surface consists of the 3F surface and a very weakly adsorbed F2 molecule. AlF3 is a strongly ionic material; consequently, the dominant theme governing its surface stability is stoichiometry. It is energetically favourable for a slab of AlF3 to consist

182

Functionalized Inorganic Fluorides

of exactly three times more F ions than Al ions, allowing the formation of an Al3þ and three F ions. The 3F surface of -AlF3 (0001) (2 1) satisfies this requirement. The formation of a stoichiometric surface within a (1 1) cell is not possible thus reconstruction is expected to occur. This principle simplifies the search for low energy AlF3 surfaces. The 110) and (0112) surfaces of -AlF3 are shown in Figure 6.4. These lowest energy (2 surfaces were calculated within the minimum cell size required to obtain a stoichiometric surface. It is, of course, possible that further surface reconstructions could occur; however, currently it is not feasible to consider these and so, in the absence of experimental determination of the surface periodicity, the simplest cell is adopted. (a)

Al(5)

Al(6)

(b)

– (0112) (1 × 1) SE = 0.94 Jm–2

Al(5)

– (2110) (2 × 1) SE = 1.04 Jm–2

Figure 6.4 The structure of the stoichiometric (011 2) (1 1) and (2110) (1 2) a-AlF3 surfaces. The F ions are represented by large spheres and the Al ions by small spheres

The surfaces of -AlF3 (100), (010) and (001) were also studied. Two or three different stoichiometric terminations for each surface were constructed. The relaxed structure and energy of each termination was calculated and the lowest energy surfaces are shown in Figure 6.5. Al(5) Four member F-AI-F-AI rings

row A

Al(5) row B

(100) (1 × 1) SE = 0.85 Jm–2

Al(5)

(010) (1 × 1) SE = 0.74 Jm–2

Al(4)

(001) (1 × 1) SE = 0.79 Jm–2

Figure 6.5 The structures and energies of the lowest energy b-AlF3 (100), (010) and (001) surfaces. The F ions are represented by large spheres and the Al ions by small spheres

Predictive Modelling of Aluminium Fluoride Surfaces

183

The - and -AlF3 surfaces all expose Al ions that are under coordinated, either fourfold or fivefold, compared to the sixfold coordination of the bulk. These Al ions are potential Lewis acid centres and may be responsible for the catalytic activity of AlF3. In Section 6.4 we shall characterize the properties of these sites. Analysis of the individual structures reveals that a simple model can be used to predict their surfaces energies, based on the local structure of the surface Al ions. We obtain the effective coordination number of an Al ion by counting 0.5 for each bidentate F ion and 1.0 for each monodentate F ion that it is bound to. For example, an Al ion bound to six F ions, one of which is monodentate, has an effective coordination number of 3.5 (5 0.5 þ 1). Three factors contribute to the surface energies; the surface density of fivefold and fourfold Al ions and the surface density of Al ions that do not have an effective coordination number of three. Analysis of our data shows that the surface energies (SE) can be estimated using SE ¼1=Af19:0 ðno: of fivefold AlsÞ þ 27:4 ðno: of fourfold AlsÞ þ2:6 ðno: of Al eff: coord 6¼ 3g

(6:15)

where all quantities are per surface cell (of area A). The prefactors are obtained from numerical fits to the data. The results from this analysis are shown in Table 6.1. As well as the lowest energy (100), (010) and (001) -AlF3 terminations shown in Figure 6.5, other (higher energy) stoichiometric terminations of these surfaces are also included in the table. It can be seen that this model can accurately predict the surface energies. The only large discrepancy is for the (100) T1 surface. This is most probably because the model does not take into account the distorted four-member (Al-F-Al-F-) rings that occur on this surface. The (001) R1 and R2 surface energies are over-estimated by around 10 %. These surfaces cut perpendicularly through the channels of -AlF3, hence the discrepancy may be due to the nonuniform distribution of the atoms in the surface plane. The accurate prediction of the (001) R3 surface may be due to this error being cancelled by the neglect of the distorted nature of its Al tetrahedra.

Table 6.1 The parameters used to predict the surface energies of the surface using Equation (6.15). The predicted energies and those calculated from our DFT energy calculations are displayed. Where multiple cuts of the same plane occur the different cuts have been given arbitrary labels Termination Surface ˚ 2) area (A b (100) T1 b (100) T6 b (010) S1 b (010) S2 b (001) R1 b (001) R2 b (001) R3 2) (011 (0001) 110) (2

88.5 88.5 51.1 51.1 85.7 85.7 85.7 25.7 44.0 36.6

No. of No. of fivefold Als fourfold Als 2 4 2 2 0 3 1 1 1 2

0 0 0 0 2 0 3 0 1 0

No. of Als coord 6¼ 3

Predicted energy (Jm2)

Calculated energy (Jm2)

4 0 0 4 2 6 2 2 2 0

0.55 0.86 0.74 0.95 0.70 0.85 1.24 0.94 1.17 1.04

0.85 0.86 0.74 0.95 0.79 0.93 1.24 0.94 1.18 1.04

184

Functionalized Inorganic Fluorides

Understanding the relationship between surface structure and energetics facilitates the prediction of potential surface reconstructions or alternate low energy surface termina2) surface reveals that it could reconstruct within a tions. p For p example, analysis of the (011 ( 2 2) cell to form a surface consisting of Al ions that are all bound to four bidentate F ions and a monodentate F ion. Its surface energy, predicted from our model, is 0.74 Jm–2, significantly lower than that calculated for the (1 1) cell. Ab Initio calculations confirm the prediction of the model, yielding an energy of 0.76 Jm–1, the relaxed structure of this surface is shown in Figure 6.6.

AI(5)

– (0112) (√2 × √2)

p p Figure 6.6 The structure and energy of the stoichiometric (01–12) ( 2 2) a-AlF3 surface. The F ions are represented by large spheres and the Al ions by small spheres

The morphology of a nanocrystal is determined by the minimization of its overall 1} surfaces are surface energy. To a very good approximation the {0001} and {101 essentially identical, hence we shall assume that their surface energies are identical. 110} and {1014} surfaces are also almost identical to one another. The Similarly, the {2 predicted equilibrium morphology of a -AlF3 nanocrystal, based on our calculations of its low index surfaces, is shown in Figure 6.7a. The surface area of the crystallite is predicted 2) surface. X-ray diffraction studies confirm that this is the to be dominated by the (011 predominantly exposed surface of -AlF3 crystallites. The bulk unit cell of -AlF3 has Cmcm symmetry. However, it is very close to the higher symmetry P6322 group. Therefore, to a very good approximation the (100) plane is equivalent to the (130) and 30) planes of P6322 symmetry. Similarly, the (010) plane is almost identical to the (110) (1 10) planes. The energies of these surfaces have been approximated to be the same as and (1 their (100) and (010) counterparts. The predicted equilibrium morphology of a -AlF3 nanocrystal is shown in Figure 6.7b. The (010) surface is predicted to dominate. There has been no experimental confirmation of this result due to the difficulties involved in synthesizing a large enough crystalline sample. The predominant surface exposed by both crystals consists of Al ions bound to four bidentate F ions and a monodentate F ion. A significant number of fourfold Al ions are also exposed on both surfaces. It is notable that only the -AlF3 surface is predicted to expose Al ions that are bound to five bidentate F ions. In the following section we characterize the reactivity of the different surfaces.

Predictive Modelling of Aluminium Fluoride Surfaces

185

{0112} – {2110} or – {1014}

{010} {100} {001}

{0001} or – {1011} (a) α-AIF3

Figure 6.7

6.4

(b) β-AIF3

Equilibrium morphology of (a) an a-AlF3 crystallite and (b) a b-AlF3 crystallite

Characterization of AlF3 Surfaces

In the previous section we showed that the surfaces of and -AlF3 all contain Al ions that are under coordinated; these ions are likely to be Lewis acid sites [22]. The local structure around these sites differs on the different surfaces. In particular, the Al ions bind to either four or five F ions and up to two of these F ions may be monodentate. It might be expected that the fourfold coordinated Al ions would be more reactive than the fivefold Al ions [23]. However, the difference in the coordination geometries of these sites also exerts a strong influence on the reactivity of the sites. In metal oxide systems, ˚ and tetrahedral coordination generally occurs for metal ions with radii less than 0.5 A ˚ [24]. The ionic radius of F is octahedral coordination for radii between 0.5 and 0.8 A ˚ compared to 1.36 A ˚ for O2 [25], hence we can expect the coordination properties 1.33 A of metal halides to be similar to those of their corresponding metal oxides. The ionic radii ˚ [25], which is close to the boundary between tetrahedral and octahedral of Al3þ is 0.54 A coordination and consequently the formation of a tetrahedral AlF4 site in the current system is not surprising. This observation is of great significance. Previously it has been suggested that the very strong Lewis acid sites on AlF3 materials may be due to fourfold Al ions [23]. However, formation of stable tetrahedral structures suggests that such sites may be only moderately Lewis acidic. Conversely, the fivefold coordinated Al ions, being in a distorted and truncated octahedra, may be expected to show stronger Lewis acidity. There is no universal method of quantifying the strength of a Lewis acid. The general definition of a Lewis acid as an electron pair acceptor gives rise to a number of possible quantifications of acidity, none of which is entirely satisfactory. Here, the acidic sites are characterized by computing the binding energy of the Lewis base NH3. This well defined process is commonly used both experimentally and theoretically for this purpose [26–28]. The binding energy of NH3 to sites on the low energy terminations of the and surfaces are shown in Table 6.2. The under coordinated Al sites on the -AlF3 and the 2) (1 1) terminations bind NH3 more strongly than the Al sites on the other -AlF3 (011 terminations. The Al ions on these terminations are bound to five bidentate F ions. As 2) termination is not predicted to occur on discussed previously, the -AlF3 (1 1) (011 the crystallite. This type of under coordinated Al site is, therefore, only expected to occur

186

Functionalized Inorganic Fluorides

on the -AlF3 surface, where it is predicted to occur only as a minor surface phase (see Figure 6.7). This may explain why -AlF3 is significantly more catalytically active than -AlF3. Furthermore, this type of site may be present in much larger quantities on HS-AlF3, which is highly catalytically active. The binding energies of NH3 to fivefold Al ions bound to one or two monodentate F ions and fourfold Al ions bound to one monodentate F ion are all similar in magnitude. Given that the binding energy is also dependent on the overall coverage of NH3 (it decreases with increased coverage) and the formation of hydrogen bonds with nearby F ions [5], it is not possible to distinguish between the strength of these acid sites by using the NH3 binding energies alone. The binding energy of NH3 to the fourfold Al ions bound to two monodentate F ions on the -AlF3 (010) termination is small in comparison to the other under coordinated Al sites considered. Table 6.2 The binding energy of NH3 to the various a- and b-AlF3 terminations. (The binding energies are corrected for BSSE using the counterpoise scheme [8].) Termination Al ion No. of monocoordination dentate F ions (011 2)a (011 2)b (2 110) (0001) b (100) T1 b (100) T6 b (010) S1 b (001) R1 b (001) R1 a b

5 5 5 5 4 5 5 5 4 4

0 1 1 2 1 0 1 1 1 2

Density of NH3 (nm2) 3.9 3.9 5.5 2.3 2.3 2.3 4.5 3.9 1.2 1.2

Hydrogen bonds NH3 binding ˚) formed (A energy (eV) 2.05, 2.05 1.92 1.64, 2.06 1.76, 1.79 1.61 1.97 1.99 1.98 1.94

1.79 1.34 1.38 1.46 1.56 1.73 1.40 1.37 1.43 1.00

The surface predicted within a (1 cell. p 1) p The surface predicted within a ( 2 2) cell.

The interaction of a Lewis acid with a Lewis base is often described in terms of the donation of an electron pair from the base to the acid. Analysis of the electronic structure of NH3 adsorbed to an AlF3 surface suggests that this is not an accurate description of the process. The Mulliken population analysis shows that there is negligible charge transfer between the two species. There is, however, a large electrostatic potential above the exposed under coordinated Al ions. A simple model can be used to estimate the electrostatic contribution to the binding energy. A single NH3 molecule is removed from the surface and the electrostatic potential is computed at the points in space previously occupied by the N and H ions. The electrostatic contribution to the binding energy is then estimated by multiplying this potential by the charge on the N and H ions of an isolated NH3 molecule computed at the structure of the adsorbed molecule. The estimated electrostatic contribution to the binding energy of NH3 on the -AlF3 (100) termination is –2.0 eV (the overall binding energy is –1.7 eV). The electrostatic contribution is very large, suggesting that the total binding energy of the molecule to the surface is dominated by electrostatic interactions. The electrostatic energy is greater than the calculated binding

Predictive Modelling of Aluminium Fluoride Surfaces

187

energy, in part, because the repulsive forces between the NH3 molecule and the surface are not considered in this analysis. The contribution to the binding energy of the relaxation of the adsorption site and NH3 molecule is also neglected. Temperature programmed desorption (TPD) is used experimentally, to measure the strength of Lewis acid sites on materials such as AlF3. To compare our theoretical results with data from TPD experiments it is necessary to calculate how the binding energy of NH3 is affected by the surface coverage The binding energies of NH3 to the under coordinated Al ions on two low energy -AlF3 (100) terminations have been calculated as a function of NH3 coverage [5]. The terminations are labelled, for historical reasons [22], as T1 and T6. The T1 termination is the lowest energy termination that has been discussed in this chapter and is shown in Figure 6.5. The T6 termination, shown in Figure 6.8, is an alternative stoichiometric surface, its surface energy is slightly higher than that of the T6 termination at 0.86 Jm–1. The surface consists of fivefold Al ions bound to 4 bidentate and a monodentate F ion, it is similar to the low energy (010) structure, differing only in that its AlF3 motifs are rotated by approximately 30% and there is a higher density of them. Consequently, we expect the adsorption of NH3 to this surface to be representative of adsorption to the majority of sites on -AlF3. AI(5)

(100) (1 × 1) T6 SE = 0.85 Jm–2

Figure 6.8 The structure of the T6 b-AlF3 (100) surface. The F ions are represented by large spheres and the Al ions by small spheres

The NH3 binding energies, calculated as a function of the occupancy of the nearest neighbour sites, were used to parametrize a kinetic Monte Carlo model to simulate NH3 desorption from the surface, as discussed in Section 6.2.4. The resultant desorption curve is shown in Figure 6.9a. The predicted spectra contains three peaks, which can be assigned to the superposition of two spectra from each termination, each of which contains two peaks. The peak that occurs at the lower temperature from each termination is due to desorption from regions of high local coverage in which NH3 is destabilized by, direct and surface mediated, intermolecular repulsions. The peak that occurs at the higher temperature is due to low coverage desorption in which most or all neighbouring sites are unoccupied. The predicted spectrum is compared to that measured on several polycrystalline -AlF3 samples [29] in Figure 6.9b. The experimental spectra are all generally consistent with the predicted spectrum. Each consists of three distinct peaks/shoulders, at around 200 °C,

188

Functionalized Inorganic Fluorides

300 °C and 420 °C The spacing of the peaks is in excellent agreement with the prediction while the absolute temperature is slightly higher (70 °C) than that predicted. This strongly indicates that the sites on the model surface are representative of the local geometries that occur on -AlF3 samples under reaction conditions. It is notable that this data also assigns the number and the position of the peaks in the TPD spectrum to a combination of surface site and inter-molecular interactions and mitigates strongly against the simplistic interpretation of peaks in TPD spectra as indicative of particular surface sites or species.

T1 & T6 (+70C) exp. A exp. B exp. C

Intensity (AU)

Intensity (AU)

T1 T6 T1 and T6

50

100 150 200 250 300 350 400 450 500 Temperature (°C)

50

100 150 200 250 300 350 400 450 500 Temperature (°C)

(a)

(b)

Figure 6.9 The TPD curves obtained from: (a) Lattice Monte Carlo simulations parametrized from the NH3 DFT binding energies from two low energy b-AlF3 (100) (labelled T1 and T6). (b) Kemnitz et al. [29] and the results obtained from our lattice Monte Carlo simulations, shifted by 70 °C to higher temperatures

6.5

Surface Composition under Reaction Conditions

The AlF3 surfaces discussed so far have all been clean and consisted only of Al and F ions. It is, however, well known that AlF3 will hydrolyse and adsorb water to form a surface that also contains O and H ions. In Section 6.2.2 we showed how the energy of such a surface can be calculated as a function of H2O, HF and H2 chemical potentials. If the surface energy of all the structures that could conceivably occur are calculated then the lowest energy surfaces can be plotted as a function of H2O, HF and H2 chemical potential. The strong ionic character of Al and F ions implies that only stoichiometric slabs which maintain charge balance will be present in the phase diagram. OH ions must, therefore, be substituted for F ions, maintaining a stoichiometry of AlF3x(OH)x. This also implies that the surface energies will be independent of the chemical potential of hydrogen, as the prefactor to this term, given in Equation (6.7), will always be zero for such surfaces. Here, geometries were created by replacing up to three surface F ions by OH ions on each surface of interest. The energies of all possible structures in which one, two or three of the surface F ions are replaced by OH ions were calculated. Adsorption of HF and H2O above the under coordinated Al ions on each of the clean and hydroxylated surfaces were also calculated.

Predictive Modelling of Aluminium Fluoride Surfaces

6.5.1

189

The a-AlF3–x (01–12) Termination

In Section 6.3.2 we showed that the surface of -AlF3 crystallites is dominated by the 2) surface. The lowest energy termination is predicted to occur within a (p2 p2) (011 cell, a metastable termination within a (1 1) was also predicted. Hydroxylation and 2) (1 1) and the molecular p p adsorption on the structures obtained from both the (011 ( 2 2) terminations were considered. The phase diagram, combining both of these terminations, is shown in Figure 6.10. The clean (non hydroxylated) termination is labelled the 3F termination, the hydroxylated surfaces are labelled according to the number of F ions that are replaced by OH ions per surface unit cell (i.e. 2F-1OH, 1F-2OH p andp3OH). The structures that appear in this phase diagram are all derived from the ( 2 2) 3F termination, except for the 3F þ H2O and 3F þ HF terminations, which are derived from the (1 1) 3F termination. The structures of the 2F–1OH, 1F-2OH and 3OH terminations, that occur in the phase diagram, are shown in Figure 6.12. The bidentate F ions are preferentially substituted for OH ions. Furthermore, it is preferable to replace F ions at positions where the Al-F-Al angle is relatively small (140°). Replacing an F ion where the Al-F-Al angle is larger (165°) is energetically unfavourable as it results in a large distortion of the surface, to form an Al-O-Al angle of approximately 140°. p p HF and H2O bind more strongly to the (1 1) 3F termination than the ( 2 2) 3F termination (Table 6.3), consequently, the (1 1) 3F terminations with adsorbed p pHF or H2O are more stable than the corresponding structures derived from the ( 2 2) 3F termination. This observation is consistent with the larger binding energies of NH3 to the (1 1) termination and the prediction that the Al ions on this termination are more reactive. hydrogen bonding also occurs on the (1 1) termination compared to p Stronger p the ( 2 2) termination. The H2O molecule hydrogen bonds via both of its hydrogens on the (1 1) termination. After adsorption of HF, the HF bond length and its hydrogen ˚ that is, an FHF species is formed. bond with a nearby F ion are of the same length, 1.15 A This behaviour has previously been seen after adsorption of HF on -AlF3 surfaces [30]. The structures of the (1 1) 3F termination after adsorption of HF and H2O are shown in Figures 6.11a and 6.11b respectively. Adsorbed HF acts as a strong Bro¨nsted acid as it can easily give up its proton, for instance to protonate nearby OH groups. It may be that the catalysis of some reactions requires the availability of both a strong Lewis acid site and a Bro¨nsted acid site. The phase diagram in Figure 6.10 is plotted for effective chemical potentials ranging from 0.0 to –2.9 eV. The corresponding partial pressures at 300 K and 600 K are also displayed. The full range of the phase diagram, shown in Figure 6.10, is not practically accessible at any given temperature. Regions of the phase diagram that are accessible at 300 K and 600 K are marked by boxes in Figure 6.10. The lower limits for the H2O and HF partial pressures are set at 10–10 and 10–15 atm respectively. These values are estimates of typical partial pressures expected under UHV conditions. The partial pressure of HF is significantly less than that of O2 in normal atmospheric conditions, hence a value of 10–15 atm is used as an estimate of UHV conditions. The upper limits are obtained from the vapour pressure of H2O and HF. At 300 K the vapour pressures of H2O and HF are 0.036 atm and 1.3 atm respectively. The vapour pressure of H2O at 600 K is 0.12 atm. The upper HF partial pressure is limited by experimental procedures and safety concerns, a maximum pressure of 5 atm is assumed.

190

Functionalized Inorganic Fluorides H2O partial pressure at 600K (atm) –15

–10

10

1e–40

–5

10

10

10

1e–30

1e–20

0

10

1e–10

5

1e+00

H2O partial pressure at 300K (atm) 10–40

10–30

10

–20

10

–10

10

0

0

3F* + HF

3F* + H2O

–1

–1.5

3F

3OH + H2O

–2

10

–10

10

–20

10

–30

2F-1OH 1F-2OH

10

5

10

0

10

10

–5

–10

HF partial pressure at 600K (atm)

HF chemical potential (eV)

0

HF partial pressure at 300K (atm)

10

–0.5

3OH

–2.5

Dissolution 10–15

–40

10

–2.5

–2

–1.5

–1

–0.5

0

H2O chemical potential (eV)

Figure 6.10 The stable a-AlF3 (011 2) surfaces, including terminations derived from the (1 1) p p and ( 2 2) 3F terminations, as a function of HF and H2O effective chemical potential (defined in Equation (6.9)) and partial pressure and temperature. The terminations derived from the (1 1) 3F termination are denoted by an asteric. The area within the small rectangle is the accessible region of the phase diagram at 300 K and the region within the large rectangle is the accessible region at 600 K (see text for details) Table 6.3 The binding energy of H2O and HF to the various a-AlF3 3F terminations. (The binding energies are corrected for BSSE using the counterpoise scheme [8].) Termination

2) (1p 1)p (011 (011 2) ( 2 2)

H2O Binding energy (eV)

HF Binding energy (eV)

1.60 1.06

1.26 0.81

Competition between hydroxylation and fluorination leads to a number of stable phases. Relatively small changes in reaction conditions can alter the surface very significantly. At 300 K the (1 1) 3F þ H2O termination is predicted to be the thermodynamically stable surface at most relevant HF and H2O partial pressures. The reactive Al ions on this

Predictive Modelling of Aluminium Fluoride Surfaces

191

H2O

FHF–

–

(b) 3F + H2O

(a) 3F + HF

Figure 6.11 termination

The structures of (a) HF and (b) H2O adsorbed on the (011 2) (1 1) 3F

OH–

(a) 2F-1OH

OH–

(b) 1F-2OH

OH–

(c) 3OH

Figure 6.12 The structures of (a) the 2F-1OH, (b) the 1F-2OH and (c) the 3OH a-AlF3 (011 2) p p ( 2 2) terminations. The F ions are represented by large spheres and the Al ions by small spheres

p p termination are shielded by adsorbed H2O molecules. At 600 K the ( 2 2) 3F termination is expected to dominate, unless the HF partial pressure is low and the H2O partial pressure is high, under which circumstances the surface is predicted to be hydroxylated. At very low partial pressures of HF and high partial pressures of H2O the surface is unstable with respect to complete hydroxylation of the crystallite. Observations of amorphous HSAlF3 have shown that, left exposed to air over a period of several months, it will undergo a transition to a hydroxylated pyrochlore structure [31]. It is important to note that the phase diagram is only based on thermodynamic considerations. The kinetic barriers to phase transitions are not considered. Forpinstance, p there is likely to be a considerable barrier to the transition from the (1 1) to the ( 2 2) structure as it requires the cleavage and formation of p several p Al-F bonds. -AlF3 is usually synthesized at elevated temperatures, at which the ( 2 2) 3F termination will dominate. It may be that the transition to the (1 1) 3FþH2O termination upon cooling to room temperature is kinetically hindered. Conversely, catalytically active HS-AlF3 is synthesized using sol-gel methods that proceed at lower temperatures [32, 33]. Under these conditions it is speculated that structures that are similar to those found on the (1 1) 3F termination form.

Functionalized Inorganic Fluorides

192

The a-AlF3 (0001) Termination

6.5.2

The phase plot for the (0001) (1 2) surface is shown in Figure 6.13. This surface is 2) termination. The flexibility of the monodentate hydroxylated more readily than the (011 F ions on this surface facilitates the formation of hydrogen bonds, which helps to stabilize the hydroxyl groups. Hydroxylation of F ions below the surface Al ions was not considered. There are both fivefold and fourfold Al ions exposed at the (0001) surface. It is possible to adsorb up to three molecules per unit cell; two to the fourfold Al and one to the fivefold Al. This leads to a very large number of possible permutations of molecular adsorption geometries. The most stable surfaces that involve adsorbed molecules are those where the molecules form strong hydrogen bonds to surface F and OH ions. Such bonds are usually formed to monodentate ions as they have greater flexibility than bidentate F ions. The structures of the 3F terminations after adsorption of two H2O and two HF molecules H2O partial pressure at 600K (atm) 10

–15

–10

–5

10

0

10

10

10

5

H2O partial pressure at 300K (atm) –40

10–30

10

–20

10

–10

10

10

0

3F + 1HF + 2H2O

0 3F + 3HF

HF chemical potential (eV)

–1

3F + 1HF +1H2O

3F + 3H2O

3F + HF 3F + 2H2O

10

–10

10

–20

10

–30

3OH + 3H2O

–1.5

2F-1OH + 2H2O

3F

1F-2OH + 2H2O

–2 2F-1OH

3OH + 2H2O

3OH

10

5

10

0

10

–5

10

–10

10

–15

HF partial pressure at 600K (atm)

3F + 2HF

0

HF partial pressure at 300K (atm)

10

–0.5

–2.5 Dissolution

10–40 1F-2OH

–2.5

–2

–1.5

–1

–0.5

0

H2O chemical potential (eV)

Figure 6.13 The stable a-AlF3 (0001) surfaces as a function of HF and H2O effective chemical potential (defined in Equation (6.9)) and partial pressure and temperature. The area within the small rectangle is the accessible region of the phase diagram at 300 K and the region within the large rectangle is the accessible region at 600 K (see text for details)

Predictive Modelling of Aluminium Fluoride Surfaces

193

are shown in Figures 6.14b and 6.14c respectively. After adsorption of two HF molecules to the 3F termination, one of the molecules forms an FHF- species, where the two H-F bonds are of equal length, while the other forms a strong hydrogen bond. In the later case ˚ and the hydrogen bond is length 1.30 A ˚ . The (0001) surface the HF bond length is 1.05 A 2) (p2 p2) generally adsorbs HF and H2O molecules more strongly than the (011 termination, due to the formation of hydrogen bonds.

OH–

(a) 2F-1OH

HF + F–

FHF–

(b) 3F-HF

H2O

(c) 3F + H2O

Figure 6.14 The structures derived from the (0001) surface. (a) The 2F-1OH termination, (b) HF adsorbed on the 3F termination and (c) H2O adsorbed on the 3F termination. The F ions are represented by large spheres and the Al ions by small spheres

At 300 K a range of terminations can be expected, depending on the HF and H2O partial pressures. Under normal laboratory conditions, three H2O molecules are predicted to adsorb on the 3F termination. At 600 K, a large number of different terminations occur. Under typical reaction conditions, for instance 20% humidity and an HF partial pressure between 10–1 and 10–5 atm two H2O molecules are predicted to adsorb on the 3F termination. In conclusion, our study of the AlF3 surfaces as a function of HF and H2O chemical potential shows that the structure of hydroxylated surfaces is predominately determined by maximizing hydrogen bonding at the surface. At elevated temperatures the surfaces will be partially hydroxylated unless they are exposed to a fluorinating agent. The phase diagrams for the two different -AlF3 surfaces studied show many similarities to one another and hence we predict that the phase diagrams of the -AlF3 will show similar characteristics.

6.6

Characterization of Hydroxylated Surfaces

Understanding how the effectiveness of an AlF3 catalyst is dependent on the extent to which its surface is hydroxylated is important in the drive to a better understanding of, and an ability to control, its catalytic properties. In Section 6.4 we characterized the Lewis acidity of non-hydroxylated AlF3 surfaces from calculations of the binding energy of NH3. This is not, however, a reliable method for characterizing the acidity of hydroxylated surfaces. NH3 strongly interacts with OH groups and in some situations adsorbed NH3 will deprotonate OH groups to form NH4þ ions.

194

Functionalized Inorganic Fluorides

An alternative method for characterizing the strength of Lewis acid sites is to consider the stretch frequency of adsorbed CO molecules. The CO stretch frequency is blue shifted after it is adsorbed on Lewis acid sites compared to an isolated CO molecule. The greater the blue shift, the stronger the Lewis acid. Bro¨nsted acid sites can also be characterized from CO adsorption. A CO molecule can form weak hydrogen bonds to surface OH groups; this interaction blue shifts both the CO and the OH stretch frequencies. We calculated the stretch frequency of CO adsorbed to the under coordinated Al ions on the clean and hydroxylated (100) and (010) -AlF3 surfaces. The Al sites on the clean (100) surface are bound to five bidentate F ions, and have previously been shown (using NH3 binding energies as a measure) to be the strongest types of Lewis acid site on AlF3. The Al ions on the (010) surface are bound to four bidentate and a monodentate F ion, and they display weaker Lewis acidity. The observed and calculated stretch frequencies for gaseous CO are 2143 cm1 and 2220 cm1 respectively. The calculated value for gaseous CO is in line with those in the literature for the B3LYP functional [34]. As in general B3LYP overestimates stretch frequencies, our calculated CO frequencies have been multiplied by a scaling factor of 0.965 (2143/2220) to take into account the discrepancy between experiment and theory. The calculated CO stretch frequencies are shown in Table 6.4. Calculations were performed at half monolayer coverage. The labelling of the F ions that were substituted for OH ions is shown in Figure 6.15. The local geometries after adsorption of CO on a selection of these surfaces are shown in Figure 6.16.

FD FC

FB FE

FD

FE FA (a) β-AIF3 (100)

FA

FC

FB

(b) β-AIF3 (010)

Figure 6.15 The local structure of the non hydroxylated surface sites on (a) the (100) T1 termination and (b) the (010) termination after adsorption of CO. Full surface calculations were performed; these pictures just show the local structure of the adsorption site

It can be seen from Table 6.4 that the stretch frequency of CO adsorbed on under coordinated Al ions is dependent on the extent to which the local ions have been hydroxylated. Substitution of F ions for OH ions reduces the stretch frequency. The position of the substituted OH group in the AlF5-x(OH) species also effects the CO stretch frequency. When an F ion is substituted for an OH ion the resultant Al-(OH)-Al angle is always smaller than the equivalent Al-F-Al angle. It was shown previously that it is energetically more favourable to hydroxylate at positions where the Al-F-Al angle is small. If the Al-F-Al angle is small (i.e. at the FA and FB ions) then the change in the angle after OH substitution is relatively small and the change in the CO shift is correspondingly small. If the Al-F-Al

Predictive Modelling of Aluminium Fluoride Surfaces

(a) 98 cm–1

(i) 82 cm–1

(d) 78 cm–1

(j) 72 cm–1

(e) 75 cm–1

(f) 24 cm–1

(p) 29 cm–1

(g) 26 cm–1

(q) 26 cm–1

195

(h) 24 cm–1

(r) 12 cm–1

Figure 6.16 The local structure of the surface sites after adsorption of CO. The top row shows structures of the sites on the (100) surface while the bottom row shows structures that occur on the (010) surface. The labelling of the figures refers to the labels given to the structures in Table 6.4. Full surface calculations were performed; these pictures just show the local structure of the adsorption site Table 6.4 Calculated shifts in CO frequency, referring to the gas phase, for adsorption to under coordinated Al ions on the b-AlF3 (100) T1 and (010) terminations. The labelling of the F ions that are substituted for OH ions is shown diagrammatically in Figure 6.15 The b-AlF3 (100) termination Label Al environment Substituted F ions a AlF5 FA b AlF4(OH) c AlF4(OH) FB FC d AlF4(OH) FB, FC e AlF3(OH)2 FB, FC, FD f AlF2(OH)3 g AlF(OH)4 FB, FC, FD, FE FA, FB, FC, FD, FE h Al(OH)5

CO stretch freq. (cm1) 2241 2238 2231 2221 2218 2167* 2169* 2167*

Shift in freq. (cm1) 98 95 88 78 75 24 26 24

The b-AlF3 (010) termination Label Al environment Substituted F ions i AlF5 FA j AlF4(OH) FC k AlF4(OH) FE l AlF4(OH) m AlF3(OH)2 FA, FB FA, FC n AlF3(OH)2 FA, FB, FC o AlF2(OH)3 FA, FC, FD p AlF2(OH)3 q AlF(OH)4 FA, FB, FC, FD FA, FB, FC, FD, FE r Al(OH)5

CO stretch freq. (cm1) 2225 2214 2210 2188 2209 2196 2176 2172* 2169* 2155*

Shift in freq. (cm1) 82 72 67 45 66 53 33 29 26 12

* Adsorption on a Bro¨nsted site.

196

Functionalized Inorganic Fluorides

angle is close to 180° (i.e. at the FC and FD ions on both surfaces and the FE ion on the (100) surface) then the change in the angle is much greater, leading to a greater distortion of the surface, consequently, the decrease in the CO shift is much greater. If more than one of these F ions is replaced by an OH group then the distortion induced at the surface is large enough for the CO to preferentially bind via one or more OH ions (i.e. at a Bro¨nsted site) if possible. Our collaborators at the Laboratory of Catalysis and spectrometry at the ENSI (E´cole Nationale Supe´rieure d’Inge´nieurs) in Caen and the University of Caen have measured the IR spectrum for CO adsorbed to clean and hydroxylated -AlF3 [35]. This work is discussed in detail in Chapter 4. Our results are used to interpret the data obtained from these experiments, enabling a greater understanding of the structure of AlF3 surfaces than could have been obtained from theory or experiment alone. In summary, both the experimental and theoretical results show that the strength of the Lewis acid sites is significantly reduced when the surface is hydroxylated.

6.7

Surface Catalysis

In Section 6.4 several AlF3 surfaces were characterized by their interaction with NH3 and it was shown that the catalytically most active surfaces consist of Al ions bound to five bidentate F ions where alternate Al ions are also bound to a monodentate F ion. It is suggested that the active sites are the under coordinated (uncapped) Al ions. The -AlF3(100) surface contains Al ions of this type and hence it is predicted that this surface of -AlF3 acts as a catalyst for many halide exchange reactions. In the previous section we showed that hydroxylation of the surfaces reduces the reactivity of the under coordinated Al sites. It is therefore imperative that surfaces are fully fluorinated before they are used as Lewis acid catalysts. In many cases the reactants are fluorinating reagents themselves and will effectively fluorinate and hence activate the surface. Very often a newly exposed surface will not act as an effective catalyst for several minutes or even hours, during which time it is thought that the reactants are acting as fluorinating agents. -AlF3 is known to catalyse several halide exchange reactions. One of the simplest of these reactions is 2CCl2 F2 ! CCl3 F þ CClF3

(6:16)

In this section we shall attempt to understand the mechanism by which this reaction proceeds on the -AlF3(100) surface. Although such reactions are often used to characterize the catalytic properties of AlF3 surfaces, the kinetics of these reactions and the mechanisms by which they proceed are poorly understood. It is known, however, that -AlF3 does not just offer adsorption sites for the reactants but that it is directly involved in the dismutation of CCl2F2 [36–41]. The reaction is thought to proceed in a nonconcerted manner; that is, a sequence of fluorination and chlorination reactions occur at the catalyst surface [42]. We propose that the dismutation of CCl2F2 to form CClF3 and CCl3F occurs as a two step process. The first step involves the adsorption of a CCl2F2 molecule via its Cl ion to an under coordinated Al ion and the subsequent dissociation of the C-Cl bond and the

Predictive Modelling of Aluminium Fluoride Surfaces

197

formation of a C-F bond with a nearby surface F ion. The newly formed CClF3 molecule then desorbs, leaving a Cl ion at the surface. The second step of the reaction involves the adsorption of a second CCl2F2 molecule, this time via its F ion. The Cl-F bond is broken and a bond is formed between the C and the Cl ion previously left behind at the surface to form a CCl3F molecule which then desorbs from the surface. This two step reaction mechanism can be written as CCl2 F2 þ Fsurf ! CClF3 þ Clsurf

(6:17)

CCl2 F2 þ Clsurf ! CCl3 F þ Fsurf

(6:18)

We investigated the energetics of this proposed reaction pathway. We initially considered the adsorption of CCl2F2 to the -AlF3 (100) surface. We then calculated the structure and energetics of the transition barriers for reactions 6.17 and 6.18.

6.7.1

Molecular Adsorption

The -AlF3 (100) surface, shown previously in Figure 6.5, consists of two rows of surface Al ions, alternate Al ions along these rows are under coordinated. In the proceeding discussion we shall refer to the upper most row as row A and the other row as row B. ˚ gap between adjacent pairs of rows on the surface, hence, they can be There is a 9 A considered to be independent of one another. Various geometries for the adsorption of CCl2F2 on the -AlF3 (100) were considered. These included adsorption via the molecule’s F and Cl ions to Al ions on both row A and row B of -AlF3 (100). The CCl2F2 molecules were adsorbed in a number of different orientations and the largest binding energies, as a function of orientation, are shown in Table 6.5. Structures consisting of CCl2F2 adsorbed F down and Cl down on row A are shown in Figure 6.17. In the case of adsorption via the Cl ion a second adsorption geometry, in which the molecule is rotated by approximately 180° is also shown. Table 6.5 The binding energies (with and without corrections for BSSE) for CCl2F2 adsorbed at half monolayer coverage on the b-AlF3 (100) surface Adsorption ion

F F Cl Cl

Site of adsorption

A B A B

Binding Energy (eV) No BSSE

BSSE

0.19 0.18 0.13 0.14

0.08 0.08 0.03 0.03

The adsorption energies are relatively small. It was shown in Section 6.4 that the binding energy of molecules to the surface is dominated by electrostatic attraction. The charge on the halide ions of CCl2F2 is only very slightly negative, hence the molecule only binds weakly to the surface. In addition, there will be electrostatic repulsion between the surface

198

Functionalized Inorganic Fluorides

(a)

(b)

CCI2F2

(c)

CCI2F2

CCI2F2

BE = –0.19 eV

BE = –0.13 eV

BE = –0.12 eV

Figure 6.17 Adsorption of CCl2F2 on row A of the b-AlF3(100) surface. (a) Adsorption via an F ion. (b) and (c) Adsorption via a Cl ion

F ions and the adsorbed CCl2F2, neighbouring CCl2F2 molecules will also repel one another. Calculations show that CCl2F2 binds more strongly, by around 0.04 eV, when the coverage is decreased from a half monolayer to a quarter monolayer (the minimum ˚ to 9.3 A ˚ .). distance between neighbouring CCl2F2 molecules is reduced from 5.3 A The correction for BSSE is approximately 0.1 eV, although this is a typical correction for BSSE, it is a large proportion of the total calculated binding energies. This implies that the structures are unlikely to be fully optimized to their BSSE-corrected minimum energy structure. It is, therefore, difficult to estimate the true binding energy for these systems as the BSSE corrected value is likely to be an under estimate of the true binding energy. We shall therefore assume that the binding energies lie somewhere between the uncorrected and BSSE corrected values. We estimate the binding energies at low coverages, to within 0.05 eV, to be –0.20 eV when adsorbed via an F ion and –0.15 eV when adsorbed via a Cl ion.

6.7.2

Reaction Mechanisms and Barriers

Minimum energy paths and transition energies were calculated for movement along row A, along row B, from row A to row B and from row B to row A for both steps of the reaction. The results from these calculations are summarized in Tables 6.6 and 6.7. The lowest transition energies for each step of the reaction occur when the CFC molecule moves between rows. The lowest energy path for the first step of the reaction is shown in Figure 6.18. It involves the adsorption of a CCl2F2 molecule to row A (Figure 6.17c) followed by the cleavage of the C-Cl bond and the formation of a CClF3 molecule adsorbed on row B. The transition barrier associated with this mechanism is 1.48 eV. The pathway for the second step of the reaction, shown in Figure 6.19, involves the adsorption of a CCl2F2 molecule to an Al on row A on a partially chlorinated surface and the subsequent formation of a CCl3F molecule on row B. The transition barrier associated with this mechanism is 0.89 eV.

Predictive Modelling of Aluminium Fluoride Surfaces

199

The transition structures for both reactions, shown in Figures 6.18 and 6.19, are quite similar. The main difference is that one of the monodentate halide ions bound to the C atom is a Cl in one case and an F in the other case. The transition energy barrier for the first mechanism is 66% larger than the barrier for the second mechanism. The distribution of the surface F and Cl ions differs between the two transition structures. This suggests that the energy barriers may depend on the position of the neighbouring F ions. An alternative reaction pathway for the the formation of a CClF3 molecule on row A was calculated. This time all the surface F ions are initially on row A and the CCl2F2 molecule is adsorbed to row B. The energy barrier for this reaction is 1.04 eV, significantly lower than the previously calculated barrier. The surface is, however, much less likely to be in this initial state, as it requires all the F ions to be on row, which is energetically unfavourable.

Table 6.6 The resultant geometries and energetics of the transition states for the reaction CCl2F2 þ Fsurf ! CClF3 þ Clsurf. The transition state energy barrier (TS Energy) is relative to the reactants CCl2F2 ads. site

CClF3 ads. site

row A row B row A row B

row A row B row B row A

TS Energy (eV)

Al–Cl ˚) (A

Al–F ˚) (A

C–Cl ˚) (A

C–F ˚) (A

F–C–Cl angle (deg.)

1.67 1.72 1.67 1.48

2.29 2.28 2.25 2.25

1.72 1.71 1.71 1.69

2.82 3.00 3.12 3.26

2.32 2.50 2.82 3.24

118 101 81 76

1.6

Energy (eV)

1.2

0.8

0.4

0 Reaction Coordinate

Figure 6.18 The lowest energy pathway for the reaction CCl2F2 þ Fsurf ! CClF3 þ Clsurf (Equation (6.17))

200

Functionalized Inorganic Fluorides

Table 6.7 The resultant geometries and energetics of the transition states for the reaction CCl2F2 þ Clsurf ! CCl3F þ Fsurf. The transition state energy barrier (TS Energy) is relative to the reactants CCl2F2 ads. site

CCl3F ads. site

row A row B row A row B

row A row B row B row A

TS Energy (eV)

Al–Cl ˚) (A

Al–F ˚) (A

C–Cl ˚) (A

C–F ˚) (A

F–C–Cl angle (deg.)

1.09 1.10 0.89 0.99

2.27 2.28 2.24 2.22

1.71 1.70 1.71 1.71

3.21 3.13 3.61 3.61

2.72 2.82 2.84 2.87

72 70 74 78

1

Energy (eV)

0.8 0.6 0.4

0 Reaction Coordinate

Figure 6.19 The lowest energy pathway for the reaction CCl2F2 þ Clsurf ! CCl3F þ Fsurf (Equation (6.18))

6.7.3

Analysing the Kinetics of the Reaction

The small binding energies associated with adsorption of the CFC molecules imply that the overall coverage of the surface will be very low under typical conditions. The binding energy of CCl2F2, at low coverages, is around –0.15 eV for adsorption via a Cl ion (see Section 6.7.1). Using Equations 6.10 and 6.11 in Section 6.2.3 it is estimated that at 600 K approximately one in every 3 107 of the under coordinated Al ions will be covered by a CCl2F2 molecule that has adsorbed via its Cl ion. Assuming a reaction barrier, DE, of 1.48 eV the reaction rate constant (Equation 6.12) is 4 s–1. The overall turnover is hence 1 10–7 s1 per Al site. If the reaction proceeds at a site where all the local monodentate F ions are adsorbed to Al ions on row A the corresponding transition barrier to the formation of CClF3 is 1.04 eV. The reaction rate constant for this reaction is 2 104 s1.

Predictive Modelling of Aluminium Fluoride Surfaces

201

It is energetically unfavourable for the F ions to all adsorb to row A. Calculations show that only one in 1600 under coordinated Al ions on the surface will be surrounded by an under coordinated Al ion on either side of it and opposite three F ions. CCl2F2 adsorbs with a binding energy of –0.35 eV. The overall turnover for this reaction pathway is approximately 2 10–5 s1 per Al site, approximately 200 times greater than for the reaction pathway with a barrier of 1.48 eV. The second step of the reaction (Equation 6.18) involves the formation of a CCl3F molecule. The minimum energy barrier for this reaction is 0.89 eV. The turnover at 600 K is 0.03 s1 per Al site, hence it is 1500 times greater than that of the first step. The first step of the overall dismutation reaction is therefore predicted to be the time limiting step. Experimental data suggest that the turnover rate per Al site is much higher than predicted from our calculations. A typical reaction in the laboratory involves passing a flow of CCl2F2, usually mixed with helium, through a microreactor filled with the AlF3 material and analysing the gas as it exits the reactor [43]. A typical reactor might contain around 20 mg of the catalyst. The gas typically consists of a 3:1 ratio of He to CCl2F2 and a typical flow rate of the gases would be 100 cm3 min1. The residence time of the flow of gas through the reactor is around 1 s. Therefore, the number of CCl2F2 molecules passing through the reactor per second is approximately 5.6 1018 at 600 K. Assuming a catalyst of surface area of 20 mg with two under coordinated Al ions per nm (see Table 6.1) and that 4 % of these are Al ions are reactive then this corresponds to around 3 1016 active Al sites in the reactor. The turnover is therefore of the order of 100 CCl3F and 100 CClF3 molecules per second per Al site as opposed to a turnover of 2 10–5 s1 per Al site predicted from our calculations. The errors in our calculated turnovers rates are likely to be quite large as they are very sensitive to the accuracy of the transition energy barrier. For instance, a 20 % over prediction of the transition energy barrier leads to an underestimate of the turnover rate by a factor of 400. We have also shown that the rate limiting step of the reaction occurs at a defect site consisting of a row of three F ions on row A opposite three under coordinated Al ions on row B of the surface. It is hence, possible that the reaction actually proceeds via a defect site that we have not considered, such as where two adjacent F ions are opposite two under coordinated Al ions. Therefore, although our calculations predict a turnover rate that is several orders of magnitude smaller than the experimental rate, this does not necessarily mean that the proposed reaction mechanism is incorrect. In the absence of detailed experimental data the mechanism that we have suggested remains the most likely candidate for the dismutation of CCl2F2 on -AlF3. This study highlights how small changes in the AlF3 surface can have a large effect on the reaction barrier of a given reaction mechanism. This suggests that HS-AlF3 may be highly catalytically active as it contains a small number of Al ions that are in a distorted environment that significantly lowers the reaction barrier for a particular reaction mechanism.

6.8

Conclusions

The ionic properties of AlF3 determine its surface structure. The Al and F ions remain strictly Al3þ and F; hence, the surface is always stoichiometric. Under coordinated Al ions are, consequently, always exposed at the surface. These Al ions are either

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coordinated to four or five F ions. We have characterized the reactivity of the under coordinated Al ions on several surfaces via the calculation of their NH3 binding energies. Our results show that fourfold Al ions are the least reactive type of under coordinated Al ions. It is suggested that this is because the coordination geometries exert a strong influence on the reactivity of the Al ions. The fourfold Al ions form stable tetrahedral structures, while the fivefold Al ions are in a distorted and truncated octahedral environment. The most reactive surface Al ions were bound to five bidentate F ions. The surface displaying this type of Al site is not predicted to be exposed on crystalline -AlF3 samples; however, it is predicted to occur in small quantities on crystallites. It is speculated that such sites occur in higher quantities on high surface area materials. This result may explain the different reactivity of -, -and HS-AlF3. 2) and (0001) terminations of -AlF3 were The structure and composition of the (011 calculated as a function of HF and H2O chemical potentials. The phase diagrams for the two surfaces showed many similarities. Under standard atmospheric conditions the surfaces were predicted to adsorb water above under coordinated Al ions. To expose the under coordinated Al ions it was shown that the surfaces must be heated up and put under conditions of low H2O partial pressures and high HF partial pressures. Although phase diagrams for the -AlF3 surfaces were not calculated, it is predicted that they would be similar to the phase diagrams of the surfaces. 2) termination phase boundaries between the The phase diagram for the (011 p contains p structures derived from the (1 1) and the ( 2 2) surfaces. The (1 1) 3F termination consists of very strong Lewis acid sites, however it is only thermodynamically stable when its Lewis acid sites are saturated by HF or H2O. This suggests that to obtain catalytically active AlF3 it is necessary to desorb these molecules p p at a temperature below that at which the surface reconstructs to form the inactive ( 2 2) phase. The solgel process used to obtain catalytically active HS-AlF3 satisfies this condition [32, 33]. The strength of Lewis acid sites on clean and hydroxylated AlF3 surfaces were characterized from calculations of CO adsorption. This study supported the results obtained from calculations of NH3 adsorption; the strongest Lewis acid sites consist of Al ions bound to five bidentate F ions. We show, furthermore, that partial hydroxylation of the surface significantly weakens the Lewis acidity of the under coordinated Al ions. The strength of the strongest type of Al site is reduced to the strength of the majority of sites found on crystalline -and -AlF3 surfaces. Analysis of the adsorption of NH3 to the surfaces of AlF3 reveals that the binding energy is predominately due to the interaction of the molecules with the large electrostatic potential above the under coordinated Al ions. Lewis acidity is commonly associated with the donation of electrons from the base to the acid. This result shows that this is not always an accurate description of the interaction. Our detailed understanding of AlF3 surfaces finally enabled us to propose a mechanism by which the dismutation of CCl2F2 occurs on -AlF3. Our study of AlF3 materials from ab initio calculations has made significant progress towards development of a new conceptual framework for understanding the chemical reactivity of high surface area AlF3 structures. This conceptual framework will underpin efforts to design better AlF3 based catalysts. The work also provides a firm basis for the investigation of a wide variety of other ionic catalysts.

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Acknowledgements We would like to thank the EU for support of the part of this work through the 6 Framework Programme (FUNFLUOS, Contract No. NMP3-CT-2004-5005575). The calculations were performed in part on the STFC’s SCARF and NW-Grid systems, in part on resources provided by the High Performance Computing Service at Imperial College and in part on the HPCx system where computer time has been provided via our membership of the UK’s HPC Material Chemistry Consortium and funded by the EPSRC (portfolio grant EP/D504872).

References [1] R. Dovesi, V. R. Saunders, C. Roetti, R. Orlando, C. M. Zicovich-Wilson, F. Pascale, B. Civalleri, K. Doll, N. M. Harrison, I. J. Bush, P. D’Arco and M. Llunell. CRYSTAL 2006 Users’s Manual. University of Torino, 2007. [2] A. D. Becke. J. Chem. Phys., 98: 1372, 1993. [3] J. Muscat, A. Wander and N. M. Harrison. Chem. Phys. Letts., 342: 397, 2001. [4] A. Wander, C. L. Bailey, S. Mukhopadhyay, B. G. Searle and N. M. Harrison. J. Phys. Chem. C, 112: 6515, 2008. [5] C. L. Bailey, A. Wander, S. Mukhopadhyay, B. G. Searle and N. M. Harrison. J. Chem. Phys., 128: 224703, 2008. [6] C.L. Bailey, S. Mukhopadhyay, A. Wander, B.G. Searle, and N.M. Harrison. J. Phys. Chem. A 113, 4976 (2009). [7] P. Spellucci. Available from http://plato.asu.edu/sub/nlounres.html. University of Darmstadt. [8] S. F. Boys and F. Bernardi. Mol. Phys., 19: 553, 1970. [9] F. Pascale, C. M. Zicovich-Wilson, F. L. Gejo, B. Civalleri, R. Orlando and R. Dovesi. J. Comput. Chem., 25: 888, 2004. [10] G. Henkelman and H. Jo´nsson. J. Chem. Phys., 113: 9978, 2000. [11] G. Henkelman, B.P. Uberuaga and H. Jo´nsson. J. Chem. Phys., 113: 9901, 2000. [12] I. Batyrev, A. Alavi and M.W. Finnis. Faraday Discuss., 114: 33, 1999. [13] X. G. Wang, W. Weiss, S. K. Shaikhutdinov, M. Ritter, M. Petersen, F. Wagner, R. Schlogl and M. Scheffler. Phys. Rev. Lett., 81: 1038, 1998. [14] K. Reuter and M. Scheffler. Phys. Rev. B, 65: 03506, 2002. [15] K. Reuter and M. Scheffler. Phys. Rev. B, 68: 04507, 2003. [16] C. L. Bailey, A. Wander, S. Mukhopadhyay, B. G. Searle and N. M. Harrison. DL Technical Report, DL-TR-2007-004, available from http://epubs.cclrc.ac.uk/work-details?w¼40553, 2007. [17] J. W. Cahn. Interfacial Segregation. American Society for Metals, Metals Park, Ohio, 1977. [18] M. W. Chase, C. A. Davies, J. R. Downey, D. J. Frurip, R. A. McDonald and A. N. Syverud. J. Phys. Chem. Ref. Data, 14, 1985. [19] A. Wander, C. L. Bailey, S. Mukhopadhyay, B. G. Searle and N. M. Harrison. J. Mater. Chem., 16: 1906, 2006. [20] R. Hoppe and D. Kissel. J. Fluorine Chem., 24: 327, 1984. [21] A. Le Bail, C. Jacoboni, M. Leblan, R. De Pape, H. Duroy and J. F. Fourquet. J. Solid State Chem., 77: 96, 1988. [22] A. Wander, C. L. Bailey, B. G. Searle, S. Mukhopadhyay and N. M. Harrison. Phys. Chem. Chem. Phys., 7: 3989, 2005. [23] S. Chaudhuri, P. Chupas, B. J. Morgan, P. A. Madden and C. P. Grey. Phys. Chem. Chem. Phys., 8: 5045, 2006. [24] P. A. Cox. Transition Metal Oxides. Claredon Press, Oxford, UK, 1995.

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[25] D. R. Lide, editor. CRC Handbook of Chemistry and Physics 73rd Edition. CRC Press, Boca Raton, Florida, 1992. [26] L. Benco, T. Bucˇko, J. Hafner and H. Toulhoat. J. Phys. Chem. B, 108: 13656, 2004. [27] M. Elanany, M. Koyama, M. Kubo, E. Broclawik and A. Miyamoto. Applied Surf. Sci., 246: 96, 2005. [28] A. A. Lamberov, A. M. Kuznetsov, M. S. Shapnik, A. N. Masliy, S. V. Borisevich, R. G. Romanova, and S. R. Egorova. J. Mol. Catal. A, 158: 481, 2000. [29] J. K. Murthy, U. Gro, S. Ru¨diger, V. V. Rao, V. V. Kumar, A. Wander, C. L. Bailey, N. M. Harrison and E. Kemnitz. J. Phys. Chem. B, 110: 8314, 2006. [30] C. L. Bailey, A. Wander, S. Mukhopadhyay, B. G. Searle and N. M. Harrison. Phys. Chem. Chem. Phys., 10: 2918, 2008. [31] A. Makarowicz, C. L. Bailey, N. Weiher, E. Kemnitz, S. L. M. Schroeder, S. Mukhopadhyay, A. Wander, B. G. Searle and N. M. Harrison, Phys. Chem. Chem. Phys. 11 5664 (2009). [32] S. K. Ru¨diger, U. Groß, M. Fiest, H. A. Prescott, S. C. Shekar, S. I. Troyanov and E. Kemnitz. J. Mat. Chem., 15: 588, 2005. [33] E. Kemnitz, U. Groß, S. Ru¨diger and S. C. Shekar. Angew. Chem., 115: 4383, 2003. [34] http://srdata.nist.gov/cccbdb/. NIST Computational Chemistry Comparison and Benchmark Database. [35] T. Krahl, A. Vimont, M. Daturi and E. Kemnitz. J. Phys. Chem. C, 111: 18317, 2007. [36] E. Kemnitz and D. Menz. Z. Anorg. Allg. Chem., 589: 228, 1990. [37] L. Rowley, G. Webb, J. M. Winfield and A. McCulloch. Applied Catal., 52: 69, 1989. [38] G. Kijowski, G. Webb and J. M. Winfield. J. Fluorine Chem., 27: 213, 1985. [39] G. Kijowski, G. Webb and J. M. Winfield. Applied Catal., 27: 181, 1986. [40] L. Rowley, J. Thomson, G. Webb, A. McCulloch and J. M. Winfield. Applied Catal., 79: 89, 1991. [41] A. Hess and E. Kemnitz. Applied. Catal. A, 82: 247, 1992. [42] A. Hammoudeh, S. S. Mahmoud and S. Gharaibeh. Applied Catal., 243: 147, 2003. [43] E. K. L. Y. Hajime, J. L. Delattre and A. M. Stacy. Chem. Mater., 19: 894, 2007.

7 Inorganic Fluoride Materials from Solvay Fluor and their Industrial Applications Placido Garcia Juan, Hans-Walter Swidersky, Thomas Schwarze and Johannes Eicher Solvay Fluor GmbH, Hans-Bo¨ckler-Allee 20, 30173 Hannover, Germany

7.1

Introduction

Solvay Fluor produces a broad range of inorganic fluorinated products in several plants around the world. They have numerous applications [1] and this chapter can give only an overview about today’s most important industrial applications of the inorganic fluorides and elemental fluorine produced by Solvay Fluor. Grinding, brazing and welding are almost impossible without fluorides and they are also crucial to the glass and aluminium industries. In addition the use of fluorinated gases for etching, cleaning and surface treatment has become common praxis in the electronics and plastics industries during recent decades. Since NOCOLOK, which is a mixture of different potassium aluminofluorides, is of great importance for the automobile industry, it will be discussed in a more detailed way.

7.2

Hydrogen Fluoride

Hydrogen fluoride is used worldwide and is distributed as pure anhydrous HF and diluted aqueous HF.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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7.2.1

Functionalized Inorganic Fluorides

Anhydrous Hydrogen Fluoride, AHF

Anhydrous hydrogen fluoride, AHF is produced worldwide in amounts of more than 1000 kt annually by reaction of anhydrous sulfuric acid with fluorspar. Processes based on the use of fluosilicic acid from phosphate rock have not yet gone beyond the development stage [7]. AHF is a colourless liquid, with a boiling point at 19.5 C at 1013 bar pressure. It is characterized by special physical properties resulting from the association of HF molecules by means of hydrogen bonds. It is also characterized by chemical reactivity and toxicity, which make it dangerous. The greatest use of AHF is in the production of fluorinated carbon compounds, the socalled HCFCs and HFCs. These materials are essential cooling agents, foam blowing agents, fire-fighting agents, solvents and raw materials for the production of fluorinated monomers for the plastics industry. Additional major uses of AHF in organic chemistry are for Balz-Schiemann reactions [2], alkylation reactions in the petrochemical industry [1], reactions with SF4 in AHF to generate CF3- or CF2H-groups [3] and electro fluorination reactions [4]. Since the production of elemental fluorine is only possible by anodic oxidation of fluoride ions, the electrolysis of HF mixed with KF as an electrolyte is the most important electrochemical application of AHF [5]. See below for additional information regarding the application of fluorine and its downstream product sulfur hexafluoride, SF6. The separation of uranium isotopes by the nuclear industry also needs anhydrous hydrogen fluoride in the first step to transform uranium oxide into UF4. For the further oxidation of UF4 to the gaseous UF6, which is centrifuged into isotopic pure fractions in a following step, elemental fluorine is needed.

7.2.2

Hydrofluoric Acid

Anhydrous hydrogen fluoride has a very good solubility in water [8]. The resulting hydrofluoric acid is transported in drums, tank cars and rubber-lined railway tank wagons. Typical HF concentrations in water for technical applications are between 50 % and 83 %. Hydrofluoric acid is used to clean cast metals [9], copper and brass [10]. It removes efflorescence from brick and stone and sand particles from metallic castings. Frosted daily products like clouded electric bulbs [11], [12] or other etched glass like polished crystal glass [13], [14], [15] and enamels [20] are well known to everybody. For the etching of silicon in photovoltaic and semiconductor industry [16], [17], [18] especially purified e-grade material is needed. Of minor importance is the decomposition of cellulose [19], the galvanizing of iron, the etching of porosity into ceramics [21], [22], specialty metal manufacture [24] as well as the use in laundry mixtures to substitute oxalic acid, in oil well acid stimulation [25], [26] and in the flotation of ores [27]. The manufacture of metal fluorides [23] by fluoride precipitation reactions between HF and metal oxides, -hydroxides or -carbonates is of great importance for Solvay Fluor. In

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Solvay Fluor’s Bad Wimpfen plant the KHF2 and NaF processes consume hydrofluoric acid in amounts of several hundred tons per month.

7.3

Elemental Fluorine, F2

Fluorine is an almost colourless gas. In dense concentrations and in the liquid state it is seen to have a pale greenish yellow colour. Fluorine or its mixtures with N2 or Ar are transported in steel bottles, bundles of steel bottles and tube trailers. Fluorine has a characteristic odour, similar to a mixture of ozone and chlorine. As the most reactive of all the elements and the most powerful known oxidizing agent, fluorine is able to react with almost all elements and compounds, with the exception of lighter noble gases and fluorides of the highest valence. These reactions occur very often even at room temperature, sometimes explosively and often accompanied by combustion. [28, 29, 30].

7.3.1

Fluorination of Plastic Fuel Tanks

Polyethylene, the plastic employed for the production of fuel tanks, has the disadvantage of being relatively permeable to fuel. This leads to fuel losses to the extent of about 20 g per day from an average fuel tank. The automobile industry today accepts a leakage rate of approximately 2 g per day but is striving to attain a maximum level of 0.2 g per day as the standard for the future [31]. Such a figure is achievable through the process of direct fluorination. The permeation of fuel through polyethylene fuel-container walls can be reduced dramatically by the chemical process of forming a fluoride layer during an inline process called blow molding [32, 33, 34, 35] or an offline process called finishing.

7.3.2

Finishing of Plastic Surfaces

The direct surface fluorination of polyethylene components also results in a reduction of gas permeation rates for gases such as oxygen, nitrogen, carbon dioxide, sulfur dioxide, C2H2 and halogenated hydrocarbons. The permeability rate of oxygen through treated PE pipes is reduced to less than 0.5 % of the original value. At the same time both the thermal stability and the chemical resistance to acids and alkalis are improved. Fluorinated PE sheets exhibit a lowered coefficient of friction and acceptance of printing inks is improved. The finishing of plastic surfaces by direct fluorination is not limited to polyethylene. Principally, any kind of plastic or rubber surface could be fluorinated. The finishing of plastic surfaces is of great significance for the plastic foil industry since for the production processes of gluing, coating, laminating, painting and printing good adhesion is absolutely essential. These abilities can be achieved by the pre treatment with elemental Fluorine.

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Functionalized Inorganic Fluorides

F2 Mixtures as CVD-chamber Cleaning Gas

F2 mixtures with nitrogen and/or Argon offer a market available, environmentally friendly alternative to the greenhouse active gases NF3, CF4, CHF3 and SF6, which are used beside perfluorocarbons as standard material for the time being. The plasma ignitability of F2/N2 and F2/Ar mixtures was studied. Both kind of mixtures ignited and showed high etching rates towards typical metals like silicon, tungsten, tantalum and titanium. Etching trials investigating the activity towards silicon oxide, -nitrides and -oxo nitrides resulted in high etching rates at low gas consumption and good particle performance [6]. The mixtures turned out to be ideal for multi-film type cleaning processes and no major issues or blocking points for industrial use could be identified.

7.4

Iodine Pentafluoride, IF5

Iodine pentafluoride is a colourless liquid with a pungent odour. The highly reactive IF5 (melting point 9.4 C) is distillable at 100.5 C without any decomposition. Decomposition starts at temperatures above 400 C. In contact with water, iodine pentafluoride hydrolyses spontaneously forming hydrogen fluoride and iodic acid. Iodine pentafluoride is not soluble in mineral acids such as sulfuric acid and nitric acid. Only a slow, regular formation of hydrogen fluoride takes place at the interface between the liquids. Iodine pentafluoride is soluble in some fluorinated solvents such as CCl3F or higher fluorinated hydrocarbons. In accordance with its conductivity IF5 is an ionizing solvent. After reaction with Lewis acids or Lewis bases the corresponding salts are formed. Scott and Bunnett [36] have shown that IF5 is partially soluble in the organic Lewis-base 1,4-Dioxane, and that the addition of an excess of IF5 results in the precipitation of colourless dioxanate crystals. With various other electron donors (e.g. pyridine) similar 1:1 molecular complexes are also formed. Organic compounds rich in hydrogen ignite spontaneously in contact with IF5 and yield hydrogen fluoride. The reactions with organic chloro-compounds are characterized by the separation of Iodine. In the case of CCl4 the main product is CCl3 F accompanied by small amounts of CCl2F2. The reaction of carbon tetraiodide with IF5 leads to the higher fluorinated CF3 I and to small amounts of CF4. Under certain conditions IF5 reacts as a strong fluorinating agent like in the oxidation of coke to CF4 [37]. The main application of IF5 is in the production of fluoroalkyl iodides [38]. These are valuable intermediates in the synthesis of perfluoralkylated organic compounds used as fluororsurfactans and textile finish [39]. 5CF2 ¼ CF2 þ 2I2 þ IF5 ! 5C2 F5 I C2 F5 I þ n CF2 ¼ CF2 ! C2 F5 ðCF2 CF2 Þ n I IF5 is also sometimes used for etching processes in the semiconductor industry.

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7.5 7.5.1

209

Sulfur Hexafluoride, SF6 SF6 as Insulating Gas for Electrical Equipment

Solvay’s sulfur hexafluoride is a nontoxic, inert, insulating and cooling gas of high dielectric strength and thermal stability [40]. For more than 40 years the gas has been used for high and medium voltage applications. It is particularly suitable for application in both high-voltage and medium-high voltage power circuit breakers (Gas Insulated Switch Gear: GIS) as well as in high-voltage cables (Gas Insulated Lines: GIL), transformers, transducers, particle accelerators (insulation of Van der Graaf generator and high-voltage GIS), X-ray equipment and UHF transmission systems and with 99.999 % purity as an etching and chamber cleaning gas in the semiconductor and photovoltaic industry. The construction of new equipment with higher capacity and improved performance has been made possible by the excellent electrical, thermal and chemical properties of SF6. Changing from conventional dielectrics to sulfur hexafluoride, which is a nonflammable chemically inactive and nontoxic gas with high density, results in considerable space and weight savings. Improvements in the operational safety of converted equipment can also be achieved [41, 42]. Electrical discharges appearing during switching processes or generated by fault electric arcs lead to the formation of gaseous decomposition products and dusty metal compounds. Gaseous decomposition products of SF6 exhibit very characteristic warning signs even at low concentrations. These warning signs are for example pungent or unpleasant odours like ‘‘rotten eggs’’, or irritation of nose, mouth and eyes. Such irritation occurs within seconds, well in advance of any danger arising from poisoning. When handling such contaminated SF6 care must be taken not to breathe in gaseous or dusty decomposition products. In case this cannot be achieved by technical safety measures like good ventilation, personal protective equipment must be worn. This personal protective equipment needs to include items of protection for the eyes, body and breathing. Due to the ecological and economical aspects, any SF6 that is released during maintenance work or when electrical systems are decommissioned should be contained and returned to the economic cycle. This limits emissions of SF6 into the atmosphere to the extremely low leakage rate of the systems. Solvay Fluor GmbH, as a leading producer of SF6, developed the SF6 – ReUse concept at the beginning of the 1990s in order to provide SF6 users with virtually unlimited possibilities to reuse SF6 and, consequently, to eliminate emissions as far as possible.

7.5.2

SF6 Applications in Metallurgy

SF6/ inert gas mixtures are used in the production of magnesium and magnesium alloys and the cleaning of aluminium melts. It functions as a protective gas to prevent ignition, oxidation and nitride formation as well as for removing oxides and solid inclusions.

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7.5.2.1

Analytical and Medical Applications

Even at the very lowest concentrations, sulfur hexafluoride can be detected by halogen leak detectors. Sulfur hexafluoride is therefore useful as an additive to other gases as a tracer compound in leak detection. Uses as a constituent of air for meteorological measurements have also been described. The property of SF6 as a high-density gas with no toxic behaviour enables medical technologists to use it as a contrast agent in ultrasonic examinations.

7.6

Ammonium Bifluoride, NH4HF2

Ammonium bifluoride can be synthesized either by reaction of ammonium fluoride or ammonia with HF: NH4 F þ HF ! NH4 HF NH3 þ 2 HF ! NH4 HF Ammonium bifluoride (Ammonium hydrogen fluoride) is a non-metallic, white crystalline solid fluoride (m.p.: 125.6 C) with good solubility in water (~620 g/l at 20 C) [43, 44]. It is used as a fluoride source and has similar applications compared to ammonium fluoride but delivers twice as much fluoride. One characteristic of ammonium bifluoride solutions is a fairly constant, moderate acidic pH (3.2–4) in a wide range of concentrations. This is due to its high dissociation grade and the pKa of hydrofluoric acid and NH4þ (3.2 and 9.25 respectively). The acidity of the solutions and the presence of fluoride make it a good mild HF-like substance that is easy to handle and, in general, it can be applied where a reaction with an oxide is needed. Examples are i) glass etching (frosted glass) in which the reaction of ammonium bifluoride, (NH4)HF2, with SiO2 results in the formation of SiF4, NH4F and H2O, and ii) metal surface treatment or chemical polishing prior a phosphate conversion coating [45, 46]. Ammonium bifluoride is also chosen for industrial cleaning of parts that are polluted with rust, scale or other silicates [47]. Ammonium bifluoride is employed in soil acidification in order to dissolve SiO2 from oil perforations [48]. It can also be used as chemical reagent for mild fluorination reactions and as starting material for the electro chemical synthesis of NF3 [49]. Its parent compound, ammonium fluoride (NH4F), can have similar applications with the difference of the pH-value of the solutions (around pH 5) and lower fluoride deliverability.

7.7

Potassium Fluorometalates, KZnF3 and K2SiF6

Potassium fluorometalates and in particular fluorozincate and fluorosilicate have applications in the field of aluminium brazing as so-called ‘reactive’ fluxes. The nomination ‘reactive’ stands for the redox reaction that takes place between the flux (KZnF3 or K2SiF6) and the aluminium surface to be brazed. These reactions give rise to formation of elemental Zn or Si and potassium fluoroaluminate, which acts as the actual flux. The elements formed in this way combine and alloy with aluminium, diffusing through the

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base material so as to reach the corresponding eutectic mixtures, lowering locally the melting point of the aluminium and thus acting as metal fillet or clad, which forms the joint: 3 K2 SiF6 þ 4 Al

! 3 KAlF4 þ K3 AlF6 þ 3 Si

6 KZnF3 þ 4 Al

! 3 KAlF4 þ K3 AlF6 þ 6 Zn

This is the case, for instance, for K2SiF6, which in an aluminothermic reaction provides in situ the Al-Si clad (molten filler metal) normally used for controlled atmosphere brazing (CAB) of aluminium. In the case of zinc, the primary objective is not only to form a clad but to obtain a homogeneous diffusion layer into the aluminium, which is needed for some applications and which would have to be otherwise applied by costly intensive methods like electrochemical galvanization, plasma coating or thermal spray coating.

7.8

Cryolite and Related Hexafluoroaluminates, Na3AlF6, Li3AlF6, K3AlF6

Synthetic cryolite (Na3AlF6; fp: 1000 C; insoluble in water) is a very important fluoroaluminate with regard to its industrial applications. Probably the most important and best known application is the use for production and refining of aluminium in combination with other fluorides, the so called Hall-He´roult process. Cryolite lowers the melting point of the fused salts [50] and dissolves aluminium oxide to allow its use for aluminium production. Cryolite is also employed in grinding applications as an abrasive aid in the phenolic resin polymeric matrix. [51]. The industrial synthesis of cryolite is based on the reaction of an acidic fluoride source like hydrofluoric acid or fluosilicic acid with aluminium hydroxide and a sodium salt. The most common sodium source is brine: 2 H2 SiF6 þ 3 Alð OH Þ 3 ! 3 AlF3 þ 3 HF þ 2 SiO2 # þ 5 H2 O AlF3 þ 3 NaX þ 3 HF ! Na3 AlF6 þ 3 HX ð X : OH ; Cl ; HCO3 . . .Þ Other uses with smaller volumes include solid lubricant for brakes in heavy-duty application, production of welding agents, pyrotechnics and metal surface treatment. Compounds related to Cryolite are Li3AlF6 and K3AlF6. Their applications are in the field of abrasives, welding, soldering and pyrotechnics. Li3AlF6 (Fp: 790 C; solubility in water: 1 g/l) is used as a component of the melt for electrolytic aluminium [52] production in order to increase the conductivity and lower the melting point. The presence of lithium in the melt also lowers the vapour pressure of fluoride salts. K3AlF6 (Fp.: > 1000 C; partial decomposition in water into KF and K2AlF5) is commonly used as a flux component for welding and brazing and as a filler in the production of abrasive articles. Such fillers have the function of absorbing the formed heat and to prevent the resin from melting [53].

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7.9

Potassium Fluoroborate, KBF4

Potassium fluoroborate is produced by reaction of fluoboric acid (HBF4) with a potassium base like a hydroxide or a carbonate. Potassium fluoroborate (decomposes at 350 C, water-solubility: 4 g/l at 20 C) is used in the aluminium metallurgy, namely in the production of Al-Ti-B master alloy as grain refiner [54]. KBF4 also finds use in the production of abrasive wheels and soldering agents, ceramics as well as a raw material in the chemical industry. The introduction of KBF4 in grinding wheels lowers the process temperature and thus helps to reduce metallurgical burn. Fluoroborates have a special place in electroplating, which is the coating of a metallic or a non-conductive substrate with another metal. The standard mixtures for electroplating baths include: (i) (ii) (iii) (iv) (v)

A metal salt (e.g. BF4) A conducting salt (e.g. SiF62) Buffering agents to buffer the pH in the cathode area (e.g. HBF4) Porosity inhibitors (wetting agents such as sulfonates) Different additives for brightness and hardening

The beneficial effect of fluoroborates lies in the efficiency at higher current densities which is better than that of sulfate electrolytes. The preparation of the appropriate metal salts starts with the dissolution of metal oxides, -hydroxides or -carbonates in fluoboric acid. The quantities vary with bath type and metal used, between 100–400 g/l of fluoroborates depending on the application [55]. KBF4 has been also used as a conducting salt in nonaqueous electrolytes for secondary lithium-ion batteries [56]. KBF4 is an additive in boronizing baths, which are used for the surface hardening of tools by [57] providing a boride layer on carbon steel.

7.10

Fluoboric Acid, HBF4

As described in the previous section, Fluoboric acid which can be found either as an aqueous solution or dissolved in a polar organic solvent like ether is the main source for the synthesis of tetrafluoroborate salts. The acid can be produced straightforward by the reaction of boric acid with anhydrous HF [58]. H3 BO3 þ 4 HF ! HBF4 þ3 H2 O Fluoboric acid is employed as a chemical reagent and catalyst in different organic syntheses like alkyne-carbonyl coupling [59], Biginelli reactions [60] and BalzSchiemann reactions. Fluoboric acid is also used in printed circuit production for tin-lead plating [61].

Inorganic Fluoride Materials from Solvay Fluor

7.11

213

Barium Fluoride, BaF2

Barium fluoride (Fp.: 1355 C; solubility in water: 1 g/l at 20 C) is used in aluminium metallurgy, in enamel and glazing frits and in the production of arc welding agents. Single crystal barium fluoride is employed in combination with gallium arsenide as insulator for semiconductor devices [62]. BaF2 can be synthesized by precipitation in a reaction of a water-soluble barium salt with aqueous HF.

7.12

Synthetic Calcium Fluoride, CaF2

Synthetic CaF2 is obtained by the reaction of a calcium salt or -oxide (e.g. CaCO3, CaCl2, CaSO4, CaO) with hydrofluoric or fluosilicic acid. Calcium fluoride (Fp.: 1423 C; solubility in water: 0,02 g/l at 20 C) is employed in aluminium metallurgy, brake lining, glass manufacturing, enamel [63] and glazing frits production, dental applications and in the production of welding agents. As high purity single crystals it is used in the fabrication of optical elements like 157 nm lithography and spectroscopic cell windows. Doped with lanthanides like Eu or Yb, CaF2 crystals can be used for up-conversion of light [64].

7.13

Sodium Fluoride, NaF

Sodium fluoride (Fp.: 993 C; solubility in water: 42 g/l at 20 C) has been widely used for the fluoridation of potable waters (1–2 mg/l). In a related application, NaF can also be employed as an additive in toothpastes [65] as anti-caries protection, in the production of vitreous enamels and as an ingredient of wood preservation compounds [1]. Sodium fluoride is also a reagent in organic chemistry, for instance as a mild fluorinating agent in chlorine – fluorine exchange reactions or as a HF-quencher. Like other Fluorides it has also been employed as an insecticide due to its enzyme inhibitor activity [66].

7.14

Sodium Bifluoride, NaHF2

Sodium bifluoride or -hydrogenfluoride is a white crystalline powder that decomposes upon heating. It is used as a laundry sour [67] and in the cleaning of stone and brick building faces. In nickelelectroplating systems NaHF2 is used to precipitate insoluble calcium as a fluoride [68].

214

Functionalized Inorganic Fluorides

Sodium bifluoride can also increase the corrosion resistance of magnesium alloys as a component of a surface treatment system.

7.15

Potassium Bifluoride, KHF2

Potassium bifluoride is a crystalline substance (Fp.: 239 C; solubility in water: 390 g/l at 20 C) that decomposes into KF and HF at 440 C. Since this decomposition behaviour opens a way to obtain high-purity HF it is one of the small-scale uses for KHF2. A far more important application of potassium bifluoride is its use as electrolyte in the production of elemental fluorine in electrolysis cells [69]. The feed of additional HF into the reaction reduces the operating temperature of the cell to about 90 C. The relation of HF to KF should not fall below KF2HF.

7.16

Potassium Fluoroaluminate, KAlF4

Apart from its use as a flux component for welding and soldering (see below), potassium fluoroaluminate (Fp.: 580 C; solubility in water: 2 g/l) has been reported as active filler in abrasives in order to stabilize the polymeric matrix against colour change due to the heat produced by the mechanical action [70].

7.17

Fluoroaluminate Fluxes in Aluminium Brazing

Since the late 1970s, potassium fluoroaluminates have been used as key chemicals in a technology called controlled atmosphere brazing (CAB). This process is used for producing aluminium heat exchangers for the automotive-, refrigeration- and air conditioning industries. The fluoroaluminates serve as a flux for removing metal oxides and for conditioning the substrate surfaces related to filler metal flow and joining. Solvay Fluor produces and markets fluoroaluminate fluxes under the trademark name NOCOLOK. 7.17.1

Flux Composition

Aluminium brazing fluxes are inorganic fluoride salts mainly consisting of potassium fluoroaluminates of the general formula K13AlF46. Fluxes of this type are considered intrinsically noncorrosive in that they are nonhygroscopic. They do not react with aluminium whether in solid or in molten stage and can remain on the surfaces of brazed components as a thin, tightly adherent and inert residue in most standard heat exchanger applications. As manufactured, potassium fluoroaluminate fluxes typically contain a mixture of KAlF4 and K2AlF5, where the K2AlF5 may or may not be hydrated. During the brazing

Inorganic Fluoride Materials from Solvay Fluor

215

process, the material undergoes essential physico-chemical transformations. While the chief component, KAlF4, is simply heated up, the compound K2AlF5 • H2O begins to lose its crystal water starting at 90C. When the temperature is further increased within the ranges of 90 C–150 C, and 290 C–330 C, two different crystallographic modifications of K2AlF5 are formed [71]. When the furnace temperature is raised above 490 C, K2AlF5 begins to dissociate: 2 K2 AlF5 ! KAlF4 þK3 AlF6 The exact amount of K3AlF6 necessary for a eutectic flux composition (KF/AlF3 phase diagram) is obtained from the original K2AlF5 content. It is the ratio between the total amount of KAlF4 (as manufactured þ from the dissociation of K2AlF5) and K3AlF6 that forms the basis for flux melting. In fact, flux manufacturers choose the ratio of KAlF4 and K3AlF6 based on the eutectic AlF3 KF phase diagram (see below), which was first constructed in 1932 [72] and further investigated and refined in 1966 [73].

PAIF3 = 1atm. 1200

KF + melt

Temperature °C

1000

KAIF4 + melt

τ -K3AIF6 + melt 800 β -AIF3 + melt KF + τ -K3AIF6 600 KAIF4 + β -AIF3 τ -K3AIF6 + KAIF4

400

KAIF4 + α -AIF3 KF + β -K3AIF6

β -K3AIF6 + KAIF4

200 KF + α -K3AIF6 20 KF

α -K3AIF6 + KAIF4 40

60

mole % AIF3

Figure 7.1 Melting phase diagram for the potassium fluoride (KF) – aluminium fluoride (AlF3) system

216

Functionalized Inorganic Fluorides

A potassium fluoroaluminate flux must contain more than just the pure phases – and it must contain these phases in an extremely precise ratio – in order to work effectively and to meet the conditions for eutectic formation which controls the melting point. The efficiency of a flux is characterized by a combination of factors such as melting and spreading, fillet formation and clearance filling (i.e. joint formation). Yamaguchi et al. showed that the pure flux phases, namely KAlF4 and K2AlF5 are not the most efficient in optimizing these flux characteristics [74]. In fact, they showed the flux to be highly effective when a combination of KAlF4 and K2AlF5 was used. Furthermore, the presence of K2AlF5 is necessary to achieve the proper ratio of KAlF4 and K3AlF6 at brazing temperature, the necessary condition to form the eutectic. This can be explained more thoroughly by examining the following scheme: Flux Components-Phase Transformation During Brazing Major As Manufactured

KAlF4

Minor and K2AlF5 • H2O and/or 90°C–150°C

K2AlF5 290°C–330°C

K2AlF5 ( I ) + ( II )

(I)

K2AlF5 ( II ) At Brazing Temperature

490°C–495°C and 500°C–520°C

KAlF4 + K3AlF6

Melt

7.17.2

Flux and HF

All potassium fluoroaluminate fluxes will create flux fumes, with strong evidence suggesting that the vapour primarily contains KAlF4. With regard to CAB brazing (furnace brazing) where traces of moisture are always present even at –40 C dew point (which corresponds to 106 ppm H2O at atmosphere pressure), a number of compounds can be formed in the system K – A1 – F – H – O. To our knowledge there has been no academic effort to create a thermodynamic model of this system. Thus, it is impossible to predict which compounds will and will not exist, and in what temperature or humidity regimes. This is why more than one mechanism has been proposed for the generation of HF, but no unique reaction mechanism has been identified: 3 KAlF4 þ3 H2 O ! K3 AlF6 þAl2 O3 þ 6 HF 2 KAlF4 þ3 H2 O ! 2 KF þ Al2 O3 þ 6 HF

Inorganic Fluoride Materials from Solvay Fluor

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While the evidence above points to gas-phase reactions between flux fumes and water vapour for the generation of HF, Thompson and Goad [75] also proposed that AlF3 dissolved in the flux melt is subject to hydrolysis according to: 2 AlF3 þ 3 H2 O!Al2 O3 þ 6 HF What is clear is that in all cases, HF is shown as a reaction product. As for the quantity, Field and Steward have indicated that the amount of HF formed is typically 20 ppm in the exhaust of a continuous tunnel furnace [76]. Solvay’s research work confirmed that even when flux on aluminium is heated in a bone-dry nitrogen atmosphere, a small quantity of HF is still generated [77]. A source of hydrogen must be made available for HF formation even under bone-dry conditions, and this might include the reduction of aluminium hydroxide, degassing of furnace walls, leakages or other less obvious sources. The work proved that even under ideal conditions it is virtually impossible to avoid some trace HF formation. Based on the above information, the authors strongly believe that the amount of generated hydrogen fluoride depends on several factors such as: • flux load (i.e. quantity) going through the furnace – flux loading and component throughput; • temperature profile – heating rate and time at temperature; • furnace atmosphere conditions such as nitrogen flow and dew point.

7.17.3

Flux Particle Size

The activity of the flux is related to chemistry and phase composition, not to particle size [74]. Spreading is simply a liquid phase reaction unrelated to the particle size distribution of the solid phase. Figure 7.2 and 7.3 show typical particle size diagrams for NOCOLOK flux as manufactured. Standard flux for slurry application shows an average particle size between 2–6 mm; and a special grade for electrostatic application an average between 3.5–25 mm. As particle size decreases, the total surface area of the flux does the same. This allows a higher surface area of flux to be in contact with the work piece. During heating up, there may be a more efficient energy transfer to the flux with the smaller particle size. The net result is that this would affect the kinetics of melting – how quickly the flux melts – but it does not affect the melting temperature range. Figure 7.4 shows the melting action of various particle size fluxes using differential scanning calorimetry or DSC. As expected, all particle sizes melt at precisely the same temperature. The appearance of the post-braze surfaces predominantly depends on initial flux loading and uniformity of coating. Once the flux melts, it is completely liquid. In its molten state, the flux has no particles. Once the flux is liquid, it immediately spreads out and wets the surfaces. Upon cooling and solidification, the amount of flux residue and its distribution on the surface of the work-piece is related entirely to the initial flux loading, and not particle size. Particle size distribution of a flux affects the characteristics of its slurry. A finer powder will stay suspended longer because it settles slower than a coarser product. Material with

Functionalized Inorganic Fluorides

218

1.2

100 95

Sympatec 1.1

90 85 75

1.0 Kennung PEN 376

0.9

70 0.8

65 60

0.7

55 0.6

50 45

0.5

40 35

0.4

Dichteverteilung q3*(x)

Summenverteilung Q3(x) / %

80

WINDOX

30 0.3

25 20

0.2

15 10

0.1

5 0 0.05

0.1

0.5

1

5

10

50

0 100

Partikelgröβe / μm

Figure 7.2 Typical particle size distribution for NOCOLOK flux, used for wet application

100

1.1 Sympatec 1.0

90 85 80

WINDOX 0.9

Kennung B 6136/15

Summenverteilung Q3(x) / %

75

0.8

70 65

0.7

60 55

0.6

50 0.5

45 40

0.4

35 30

Dichteverteilung q3*(x)

95

0.3

25 20

0.2

15 10

0.1

5 0 0.05

0.1

0.5

1

5

10

50

0 100

Partikelgröβe / μm

Figure 7.3 Typical particle size distribution for NOCOLOK Flux Drystatic, used for dry (electrostatic) application

Inorganic Fluoride Materials from Solvay Fluor

219

PEN356 Nocolok Granulat 460 800 µm, 05.06.2001 15:01:06 PEN356 Nocolok Granulat 460 800 µm, 3,8000 mg PEN356 Nocolok Granulat 250–460 µm, 05.06.2001 14:04:15 PEN356 Nocolok Granulat 250–460 µm, 4,9000 mg PEN356 Nocolok Granulat 120–250 µm, 05.06.2001 13:02:52 PEN356 Nocolok Granulat 120–250 µm, 4,2000 mg Methode: HR_B 80,0°C 10,0 min 80,0–500,0°C 30,00°C/min 500,0–600,0°C 5,00°C/min

N2, 80,0 ml/min N2, 80,0 ml/min N2, 80,0 ml/min

20 mW

80 0

5

100

200

10

300 15

400 20

500

520 25

540 30

560 35

580

°C

40

min

Figure 7.4 DSC scans for granulated NOCOLOK flux samples a) 120–250 mm; b) 250–460 mm; and c) 460–800 mm

larger grains seems to build up more rapidly on inside surfaces of slurry tanks and spraying equipment. Regardless of the specific particle distribution of a flux, continuous agitation is necessary to prevent settling and build-up. Regular maintenance is the only way to avoid the formation of solidified material residues.

7.17.4

Flux Melting Range

According to the KF/AlF3 phase diagram, melting of potassium fluoroaluminates eutectic compositions begins slightly above 560 C [72]. As soon as the flux begins to melt, one of the components of the flux – KAlF4 – begins progressively evaporating, with a vapour pressure determined to be 0.08 mbar at 600 C [75]. Evaporation of KAlF4 causes the flux melt to change composition, and it begins to dry out. Given enough time, it is possible for the flux melt to completely dry out before reaching the maximum peak brazing temperature. A good brazing flux only needs to be available just before filler metal melting. Table 7.1 describes what happens at brazing temperature. As soon as the flux melts, it begins to dissolve the oxide layer, and this solvating process continues until the oxide is removed, even if the filler alloy has melted. Yamaguchi et al. have shown that with an increase in the K2AlF5 content, the flux will start to melt at a lower temperature so that the flux will work at a lower temperature. However, even if KAlF4 evaporation is ignored, the same study concluded that increasing the K2AlF5 content eventually prevents the flux from spreading smoothly, and therefore affects the efficiency of the flux [74].

220

Functionalized Inorganic Fluorides

Table 7.1

Steps during an aluminium-brazing process

Surface temperature of aluminum components

Temp. Action range D

Below 560 C Between 560 and 575 C Between 575 and 585 C

Duration with a heating-ramp of 15 C/min

Uniform heat-up 15 C

Flux melting

0.6 min ¼ 36 s

Depends on brazing cycle 1 min ¼ 60 s

10 C

Flux spreading and surface wetting Brazing range of AA 4045 filler alloy Brazing range of AA 4343 filler alloy

0.4 min ¼ 24 s

0.66 min ¼ 40 s

0.8 min ¼ 48 s

1.33 min ¼ 80 s

0.8 min ¼ 48 s

1.33 min ¼ 80 s

Approximately 585–605 C

20 C

Approximately 590–610 C

20 C

7.17.5

Duration with a heating-ramp of 25 C/min

Current Status of Aluminium Brazing Technology

Since the early 1980s, controlled atmosphere brazing (CAB) with noncorrosive fluxes has evolved as the leading technology for manufacturing aluminium heat exchangers. Presently, more than 700 CAB furnaces are in operation throughout the world using the NOCOLOK process. This technology offers the benefits of a flux for successful o9xide removal, working at atmospheric pressure while avoiding the disadvantages of postbraze treatments and corrosion susceptibility. The nonhygroscopic and noncorrosive potassium fluoroaluminates are used as fluxes [76]. The process in most brazing operations includes the following steps: • component forming and assembly; • cleaning and flux application; • brazing. The process sequence in brazing operations is dependent on: • heat exchanger design; • cleaning method; • flux application method. Success or failure in CAB production relies on several factors. The starting point is a good product fit up. Parts to be joined by metallurgy must have intimate contact at some points along the joint. An adequate but not excessive quantity of filler metal must be available to fill the joints. Capillary forces pull the filler into the joints. The gap tolerance is 0.1 to 0.15 mm for nonclad components. When clad products are used, intimate contact is recommended; the clad layers will act as a gaping tool.

Inorganic Fluoride Materials from Solvay Fluor

221

Another essential for reliable brazing results is a uniform flux coating on all surfaces involved in the joint formation. The main focus for achieving this task is on the cleaning and fluxing procedure. Equally important are the furnace conditions, i.e. temperature profile, temperature uniformity and atmospheric conditions. 7.17.6

Cleaning and Flux Application

Cleaning and flux application procedures depend on product designs and personal choice. Heat exchangers are classified in two groups: (a) no internal brazing required; (b) internal brazing required. For the majority of brazed heat exchangers, assembly is followed by cleaning and subsequently by flux application. However, when internal surfaces need to be fluxed, the production sequence may follow a different order, determined by when and how the flux is applied. The purpose of cleaning is to remove fabricating oils and lubricants as well as other contaminants from the surfaces. The cleaning procedure must allow for adequate flux retention and render the surfaces suitable for brazing. Cleaning results have great influence on the brazing result, postbraze product appearance and corrosion performance. There is a wide variety of different fluxing methods including: • • • • • •

low pressure spray; flooding (cascade); dipping or submerging; high pressure (atomized) spray; brushing (paste application); dry or electrostatic.

7.17.7

Wet Flux Application

The most common flux application method in CAB is by spraying an aqueous suspension. Constantly agitated flux slurries with concentrations of approximately 5–35 % solids are pumped from tanks to fluxing booths. All aluminium surfaces involved in the brazing process are coated with the slurry, resulting in a uniform flux layer. Excess flux slurry is removed with a high-volume air blow; the excess is then collected, recycled and reused in the fluxing booth. When wet fluxing is used, the heat exchangers are dried prior to brazing in a separate dry-off oven. Product entering the brazing furnace must be completely dry from water introduced via aqueous cleaning or flux slurry coating. Aluminium oxide thickness increases with temperature, time at temperature and particularly in the presence of moisture. Oxide formation will affect clad fluidity [78]. There is ˚ at only a small drop in braze ability when the oxide thickness increases from 40 to 220 A 5 g/m2 flux load. However, there is an appreciable drop in clad fluidity as the aluminium

Functionalized Inorganic Fluorides

222

˚ at 2 g/m2 flux. A reduction in braze ability can oxide thickness increases from 40 to 100 A ˚ . Proper flux loads should even be noted as the oxide thickness is increased from 40 to 60 A be maintained since low flux load appears to be quite sensitive to even small changes in the oxide layer’s thickness. 7.17.8

Dry/Electrostatic Flux Application

Since the early 1990s, some users of CAB technology have successfully implemented dry flux application methods. Based on the principles of powder-paint technology, an alternative application technique was introduced in the brazing industry. Particularly when dry fluxing is used in combination with thermal degreasing (i.e. evaporative oils and lubricants), the objective is to completely eliminate or significantly reduce the water consumption in the process. Some operations apply the flux in an electrostatic way before they thermally degrease the units. The thin layer of residual lubricants seems to improve powder adhesion [79]. Special flux qualities with adjusted particle distribution are commercially available, and contribute to an improved electrostatic application. 7.17.9

Post Braze Flux Residue

It is generally accepted that the presence of flux residues on a heat exchanger enhances its corrosion resistance [80, 81, 82]. However, it has always been difficult to quantify the level of corrosion resistance enhancement. In corrosion testing of flat panels or coupons coated with flux residue, there is no doubt that there is a beneficial effect. With heat exchangers on the other hand, the general trend shows a longer corrosion life but factors such as uniformity of flux residue coverage and variations in flux load sometimes confuse the corrosion test data. There is no indication of interactions between flux residue and coolants, refrigerants [83], turbine oils, and polyalcylene glycol lubricants. Recent experimental work has revealed the correlation of flux residue solubility and fluoride release [84]. 7.17.10

Filler Metal Alloys

Commercial filler metals are aluminium silicon alloys containing from 6.8 % to 13 % silicon. The solidus (or the point at which melting begins) is 577 C for all Al/Si filler metal alloys. However, melting occurs in a range. Before the filler metal becomes fully liquid Table 7.2 Filler metal alloy characteristics Alloy AA

Si [85]

4343 4045 4047

6.8–8.2% 9–11% 11–13%

Start Melting

Fully Molten

Braze Range

577 C 577 C 577 C

613 C 591 C 582 C

593–610 C 588–604 C 582–600 C

Inorganic Fluoride Materials from Solvay Fluor

223

and starts to flow, more than 60% of the clad material must melt. This requires a minimum (threshold) temperature for each filler metal alloy in the brazing process. Special developments for filler metal alloys are focused on the control of fluidity and flow pattern [86, 87]. By adding specific trace quantities of alloying elements (e.g., Li, Na), brazing characteristics improve. These effects appear to be relating to reduced surface tension. Under certain conditions, brazing can be accomplished with reduced flux loads or with no flux at all [88]. In Mg-containing alloys Mg and/or MgO can react with flux forming compounds such as MgF2, KMgF3 and K2MgF4. These compounds serve to poison the flux, significantly reducing its effectiveness. The limit to the amount of Mg tolerated in furnace flux brazing is 0.5 %; while around 1 % Mg is tolerable for flame brazing. It should be noted that the brazing tolerance to Mg is the total sum of the Mg concentrations in both components to be joined: ½Mg component 1 þ ½Mg component 2 ¼ ½Mg total Improved brazing results have been reported with magnesium-containing aluminium alloys (up to 0.6 % Mg) when the flux formulation contains some caesium (i.e. approximately 2%) [89]. Caesium reacts as a buffer for Magnesium by forming CsMgF3 and Cs4Mg3F10, which reduces the flux inhibition. These compounds melt at lower temperatures and interfere less with aluminium brazing [90].

7.17.11

Flux Precoated Brazing Sheet/Components

The concept of a brazing sheet that is supplied with a flux coating is very plausible. Such material would significantly change the way heat exchangers are currently manufactured in that the flux application step would be eliminated. Several patent applications have recently been filed on this premise [91, 92, 93]. Its greatest challenge is the best way to make the flux adhere to the metal surface. Throughout the forming process of the components, uniform coverage and strong adhesion are equally important. The latter is the origin of yet another concept. Flux-coating technologies have been developed for preformed components, which are primarily adapted for heat exchangers with internal brazing required, i.e. plate evaporators [94]. Flux application with a binder system allows coating of specific surface areas with a precise flux amount. It also reduces flux fall-off during assembly. Binders used for pre-fluxing must evaporate during the process without interfering with the brazing performance or leaving any contamination on the surfaces. Exhaust treatment may need to be adjusted to the presence of binder fumes.

7.17.12

Clad-less Brazing

There are several suggestions for brazing technologies where the filler metal is generated during the brazing cycle from a coating layer on a plain aluminium sheet or on extrusion material.

224

Functionalized Inorganic Fluorides

One method involves a mixture of flux with silicon powder, the NOCOLOK Sil Flux process [95]. At brazing temperature, the silicon powder diffuses into the aluminium substrate and generates Al-Si filler metal for joint formation. Sil flux can be applied with a binder to specific component surfaces, e.g. extruded tubes. In this case, the filler metal would be supplied from the tube and a clad fin sheet is not necessary. In Composite Deposition technology (CD process) [96], the filler metal for joint formation is derived from a composite powder, a compound consisting of potassium fluoroaluminate flux and Al-Si alloy. The CD powder is ‘deposited’ selectively and accurately on heat-exchanger components prior to assembly and brazing. The common challenge for all cladless approaches is the development of a reliable coating procedure. 7.17.13

Furnace Conditions

Specifications for acceptable CAB furnaces (with the use of 5 g/m2 uniform flux load) are: • furnace atmosphere: • oxygen level below 100 ppm; • dew point below –40 C; • brazing temperature and time: • ideally a uniform 600 C – 5 C (heat exchanger surface temperature); • ideally 3 min – 0.5 min from 580 C heat up to 605 C on cool down. Many production sites work with flux loads of 3 g/m2. As flux coating goes down, furnace atmospheric conditions become more critical. Most CAB furnace manufacturers produce furnaces which routinely operate in the 20–50 ppm oxygen and dew point levels of –45 to –50 C (i.e., 92–67 ppm H2O). A low dew point will keep the formation of hydrogen fluoride to a minimum [97]. A suitable treatment for the furnace exhaust (i.e. dry scrubber) is required.

7.18

Summary

Controlled atmosphere brazing using non-corrosive potassium fluoroaluminate flux is the dominant process for making all aluminium heat exchangers. There are several factors determining the success of aluminium brazing: • • • • • •

product fit-up and assembly; component cleanliness; flux application; furnace atmosphere; brazing temperature uniformity; brazing time at temperature.

This technology provides robust and stable parameters. Over the years, a significant optimization potential was achieved in all process steps as well as in equipment and machinery.

Inorganic Fluoride Materials from Solvay Fluor

225

Developments focus on further improvements in aluminium alloy composition (i.e. higher strength, good formability and higher corrosion resistance). Research activities throughout the industry have resulted in concepts for: • new flux qualities more suitable in dry/electrostatic flux application and for furnace brazing of Mg-containing alloys (up to 0.6%); • cladless brazing technologies; • prefluxed brazing sheet/ components; • reduction of consumables (e.g. flux, water). Future developments for flux brazing technology will include manufacturing more heat exchanger types and extended product ranges, some of which are currently still produced with other processes and materials.

References [1] D. T. Meshri, Advanced Inorganic Fluorides: Synthesis, Characterization and Applications, edited by T. Nakajima, B. Zemva, A. Tressaud, 661–682, Elsevier Science S.A., Amsterdam, 2000. [2] G. Schiemann, Aromatic fluoro derivatives. XIX. The borofluoride method of preparing aromatic fluoro derivatives, J. Prakt. Chemie, 140, 97–99 (1934). [3] A. I. Burmakov, B. V. Kunshenko, L. A. Alekseeva and L. M. Yagupolskii, New Fluorinating Agents in Organic Synthesis, 197–253, Springer-Verlag, Berlin, 1989. [4] R.L. Powell and T.A. Ryan, Industrial Inorganic Chemicals and Uses, edited by R. Thompson, 208–209, Royal Society of Chemistry, Cambridge, 1995. [5] D. Naumann, Fluor und Fluorverbindungen, Spezielle Anorganische Chemie Bd. 2, 3–5, Dr. Dietrich Steinkopff Verlag, Darmstadt, 1980. [6] Solvay Fluor GmbH, Hannover, 2006. [7] J. A. Aigueperse ‘‘Hydrofluoric Acid’’, Chapter 1, Fluorine Compounds, Inorganic; Ullmann’s ‘‘Encyclopaedia of Industrial Chemistry’’; Part 11, Wiley-VCH Verlag Weinheim, 2005. [8] J. F. Gall. In: Kirk-Othmer; ‘‘Encyclopedia of Chemical Technology’’; 3rd Edition, 10, 733–753 Wiley, New York 1980. [9] C. S. Wen and Y. T. Yu, Surface characterization of cast Ti-6Al-4V in hydrofluoric-nitric pickling solutions, Surface and Coatings Technology, 176, 337–343 (2004). [10] Y. Uozumi, T. Nakajima, T. Matsumura, Y. Yoshimizu and H. Tomita, ‘‘Development of an Eco-friendly Copper Interconnect Cleaning Process’’, Paper from: International Interconnect Technology Conference, IEEE 2007 4–6, 25–27, Institute of Electrical and Electronics Engineers (IEEE) 2007. [11] P. D. H. Short, Etching of glass articles, Glass., 57, 208–211 (1980). [12] W. C. Nixxon, Jr., Glass Frosting Problems and How to Correct Them, Glass Industry, 66, 14–16, 39 (1985). [13] Y. S. Lazarev and V. A. Drozhzhinov, Chemical polishing of crystal glass with concentrated acids, J. Glass and Ceramics, 47, 433–434 (1990). [14] C. Chang, T. Abe and M. Esashi, Glass etching assisted by femtosecond pulse modification, Journal Sens. Mater., 15, 137–145 (2003). [15] M. Bu, T. Melvin, G. J. Ensell, J. S. Wilkinson and A. G. R. Evans, A new masking technology for deep glass etching and its microfluidic application, Sensors and Actuators A, 115, 476–482 (2004). [16] A. Somashekhar and S. O’Brien, Etching SiO2 Films in Aqueous 0.49% HF, J. Electrochem. Soc., 143, 2885 (1996). [17] D. M. Knotter and T. J. J. Denteneer, Etching mechanism of silicon nitride in HF-based solutions, J. Electrochem. Soc., 148(3), 43–46 (2001).

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[18] I. Ro¨ver, G. Roewer, S. Patzig, K. Bohmhammel and K. Wambach, Process controlling of the etching system HF/HNO3/H2O, Proceedings of the 20th European Photovoltaic Solar Energy Conference, EPSEC (2005) by WIP Wirtschaft und Infrastruktur GmbH & Co Planungs KG-Munich. [19] D. T. A. Lamport, G. Smith, S. Mohrick, M. C. Hawley, R. Chapman and S. Selke ‘‘Hydrogen Fluoride Saccharification’’: The Key to Ethanol From Wood, (Paper prepared for the International Conference on Energy from Biomass, Brighton, England) Chem. Abstract, 96, 169 (1982). [20] W. C. Lindemann, The electric cleaning of metals for enameling purposes, J. Amer. Ceramic Soc., 3, 252–255 (1920). [21] E. A. Vasil’eva, L. V. Morozova, A. E. Lapshin, E. N. Solov’eva and V. G. Konakov, Specific features of the synthesis of porous materials based on a magnesium–aluminum spinel, Glass Phys. and Chem., 29, 490–493 (2003). [22] R. Pen˜a-Alonso, G. D. Soraru` and R. Raj, Preparation of ultrathin-walled carbon-based nanoporous structures by etching pseudo-amorphous silicon oxycarbide ceramics, J. Amer. Ceramic Soc., 89, 2473–2480 (2006). [23] Solvay Fluor; Solvay Fluorides – Brochure, 2002. [24] N. J. Sanders, Environmentally friendly stainless steel pickling, Anti Corrosion Methods and Materials, 44, 20–25 (1997). [25] H. K. van Poolen, Comprehensive study, Oil and Gas Journal, 11, 93ff (1967). [26] N. Al-Araimi and L. Jin, A high-success-rate acid-stimulation campaign – a case history, SPE101038, Presentation at the 2006 SPE Asia Pacific Oil and Gas Conference, Adelaide, Australia, 11–13 September 2006. [27] V. V. Klyachin, New effective reagents for the flotation of feldspar, Glass and Ceramics, 29, 605–606 (1972). [28] Gmelin Handbuch der Anorganischen Chemie, Vol. 8, Fluorine, supplement volume 2, Springer-Verlag, Berlin (1980). [29] M. Grayson ‘‘Fluorine’’ In: Kirk-Othmer, Encyclopedia of Chemical Technology; 3rd Edition, 10, 630–654, Wiley, New York 1980. [30] M. Wechsberg, W Schabacher, H. Niederpru¨m, S. Schneider, V. Beyl ‘‘Fluor und Anorganische Fluorverbindungen’’, Ullmanns Encyklopa¨die der technischen Chemie, 11, 4th Edition, VCH, Weinheim, 587–595, (1975). [31] Solvay Fluor GmbH, Elemental Fluorine, brochure, Solvay Fluor, 1994. [32] Sperrschichtbildung bei Kunststoff-Hohlko¨rpern, Verein Dt. Ingenieure, VDI-Ges. Kunststofftechnik, VDI-Verlag, Du¨sseldorf, 1986. [33] P. A. B. Carstens, S. A. Marais and C. J. Thompson, Improved and novel surface fluorinated products, J. Fluorine Chem., 104, 97–107 (2000). [34] A. P. Kharitonov, R. Taege, G. Ferrier, V.V. Teplyakov, D.A. Syrtsova and G.-H. Koops, Direct fluorination — Useful tool to enhance commercial properties of polymer articles, Journal of Fluorine Chemistry, 126, Issue 2, 251–263 (2005). [35] D. Masson, US Patent 5447667. Method and device for the blow molding of hollow bodies from thermoplastic material, US Patent Issued on 5 September 1995, assignee Solvay. [36] A. Scott and J. Bunnett, ‘‘A Dioxanate of Iodine Pentafluoride’’, J. Amer. Chem. Soc., 64, 2727 (1942). [37] Patent, JP 58203924, Production of Carbon tetrafluoride, Kanto Denka Kogyo. [38] E. Kissa, Fluorinated Surfactants and Repellents, CRC Press, 2nd Edition, Chapter 2, 30–31, Cleveland Ohio (2001). [39] R. E. Banks, B. E. Smart and J. C. Tatlow, Organofluorine Chemistry: Principles and Commercial Applications, 325, Springer Verlag, Berlin (1994). [40] Solvay Fluor GmbH, Sulfur Hexafluoride, product brochure, 2004. [41] Stromversorgung unter Nutzung der SF6-Technologie, ABB, Preussen Elektra Netz, RWE Energie, Siemens, Solvay Fluor und Derivate (1999). [42] SF6-GIS-Technologie in der Energieverteilung – Mittelspannung, ABB, Areva, EnBW, E.on, RWE, Siemens, Solvay Fluor und Derivate (2003). [43] Handbook of chemistry and Physics. 52nd Edition. Ed. Robert C. Weast , CRC-Press, Cleveland Ohio (1971–1972). [44] Solvay Ammonium Bifluoride Product Data Sheet.

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[79] Y. Luo and H. Hutchins, Method for Braze Flux Application, European Patent Application EP 0931619 A1. [80] T. Nakazawa, K. Tanabe, K. Ushikubo and M. Hiraga, Performance Evaluation of Serpentine Evaporator for Automotive Air-conditioning System, SAE Paper # 840384, Warrendale, PA. [81] D. Claydon and A. Sugihara, Brazing Aluminum Automotive Heat Exchanger Assemblies Using a Non-Corrosive Flux, SAE Paper # 830021, Warrendale PA. [82] W. E. Cooke, T. E. Wright and J. A. Hirschfield, Furnace Brazing of Aluminum with a Non-Corrosive Flux, Welding Journal, 23 (1978). [83] C. Meurer, D. C. Lauzon and H. Ko¨nig, Stability of R-134a in the Presence of NOCOLOK Flux Residues, SAE Paper # 980052, Warrendale, PA (1998). [84] P. Garcia and H.-W. Swidersky, Study on the hydrolysis of fluoroaluminates in post-braze flux residues, A.F.C. Holcroft International Invitational Aluminum Brazing Seminar, Novi, October 2008, conference proceedings. [85] Registration Record of International Alloy Designations and Chemical Composition Limits for Wrought Aluminum and Wrought Aluminum Alloys, The Aluminum Association, Inc. (July 1998). [86] I. Okamoto and T. Takemoto, Brazability of aluminum using Al-Si filler alloys with different compositions and microstructures, Transactions of JWRI, 10, 35 (1981). [87] D. P. Sekulic´, Behavior of aluminum alloy micro layer during brazing, Recent Res. Devel. Heat, Mass and Momentum Transfer, 2, 121 (1999). [88] D. Childree, Fluxless brazing in a controlled atmosphere furnace with a new filler alloy containing Na, International Invitational Aluminum Brazing Seminar, Detroit, October 2000, conference proceedings. [89] J. Garcia, C. Massoulier and Ph. Faille, Brazeability of Aluminum Alloys Containing Magnesium by CAB Process Using Cesium Flux, International Invitational Aluminum Brazing Seminar, Detroit, October 2000, Conference Proceedings. [90] U. Seseke, New developments in non-corrosive fluxes for innovative brazing, First International Congress Aluminium Brazing, Du¨sseldorf, May 2000, conference proceedings. [91] R. Kilmer and J. Eye, A method of depositing flux or flux and metal onto a metal brazing substrate, International Patent Application WO 00/52228. [92] N. Sucke, Partial or Complete Coating of Aluminum and Aluminum Alloy Structural Parts With a Braze Flux, and Bonding Agent Prior to Brazing, German Patent Application DE 19859735 A1. [93] A. Wittebrood, Composite Sheet Material for Brazing, International Patent Application WO 00/64626 A1. [94] M. Kojima, F. Watanabe, A. Toukana, H. Nojiri and K. Yamasoe, Flux Composition for Brazing of Aluminum Material and Method for Brazing of Aluminum Material, European Patent Application EP 0936024 A1. [95] R. S. Timsit and B. J. Janeway, A Novel Brazing Technique for Aluminum, Welding Journal, Welding Research Supplement, 119–128 (1994). [96] J. S. Coombs, Production of Powder, International Patent Application WO 94/17941; European Patent Specification EP 0682578 B1. [97] D. Lauzon, H.-J. Belt and U. Bentrup, HF Generation in NOCOLOK Flux brazing furnaces, 1998 International Invitational Aluminum Brazing Seminar, conference proceedings.

8 New Nanostructured Fluorocompounds as UV Absorbers Alain Demourgues, Laetitia Sronek and Nicolas Penin ICMCB, CNRS, Universite´ Bordeaux 1, 87 Avenue du Dr. A. Schweitzer, 33608 Pessac Cedex, France

8.1

Introduction

Titanium dioxide (TiO2), zinc oxide (ZnO) and cerium dioxide (CeO2) exhibit optical band gaps at 3.1–3.25 eV, 3.4 eV and 3.2 eV respectively1–6 at the frontier between the UV and the visible range, corresponding to charge transfers O(2p)-M(4s/3d) (M ¼ Zn/Ti) and O(2p)-Ce(4f/5d). However, they present major drawbacks: (i) they exhibit high photocatalytic activity under UV irradiation or oxidation catalytic activity, leading to the generation of reactive oxygen species7–9 that can induce a photodegradation of the organic medium in which they are dispersed (varnish, paint, polymers, etc.);10–11 (ii) the refractive index is high in the visible range, i.e. around 2.75–2.45 for TiO212 and CeO213 compounds. As a consequence of these high refractive indices a whitening of the medium is observed limiting the use of such materials as UV absorbers – in the case of wood protection, for example. An increase in transparency in the visible light region can be achieved by reducing the particle size down to the 20–50 nm range. In recent years, numerous papers have been published concerning cerium oxide CeO2 as a possible alternative (Eg ¼ 3.2 eV, n ¼ 2.2–2.4)14. The refractive index of nanocrystalline cerium oxide films is strongly dependent on the film density.15,16 For instance, for the density of the bulk CeO2 (7.21 g.cm3), the refractive index measured by ellipsometry is around 2.3515 in the visible

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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range. However for a lower calculated density, around 6.7 g.cm3, due to the polycrystalline character of the films and the occurrence of voids, the refractive index is found at 2.12.16 Electronic density calculation coupled with electron energy loss spectroscopy measurements show also that the refractive index remains rather high, i.e. around 2.4517 in the visible range, and it is necessary to get highly divided particles or aggregates with sizes around 30 nm in order to limit scattering and reach a high transparency in the visible range.14,17,18 However, CeO2 appears with a pale yellow coloration and a slight absorption in visible range. For various industrial applications, it is relevant to keep white colour and avoid any yellowish hue. Moreover, the substitution of various cations for Ce4þ or Ti4þ has been widely investigated either to enhance or to reduce the photocatalytic or oxidation catalytic activity.19–22 Ca2þ substitution for Ce4þ ions in CeO2, led to nano-sized particles with a clear reduction of the oxidation catalytic activity of ceria.23The Ce/Ca coprecipitation process was carried out from chloride precursors at neutral or basic pH, followed by oxidation with H2O2.19 Ce1xCaxO2x compositions with particle sizes around 5–10 nm exhibit UV-shielding properties with a slight reduction of the optical band gap compared to the undoped oxide.23 Yamashita et al.19 suggested that the maximum solubility limit of CaO into CeO2 is equal to or less than 30 mol%, corresponding to the following formula Ce0.7Ca0.3O1.7. Particle sizes decrease with Ca content whereas strain parameters increase in good agreement with the occurrence of oxygen vacancies.24 From T ¼ 600 C, the strain parameters decrease up to the formation of CaO at T ¼ 800 C and the particle size strongly increases in this temperature range.24 Moreover electronic properties of Ce1xCaxO2x oxides based on O K and Ce LIII-edges of XANES spectra and calculated density of states24 reveal the reduction of 2p (O) - 4f (Ce) electronic transition energy with Ca doping with a cation charge redistribution where the Ca atoms become more electropositive than in CaO while the Ce atoms are less electropositive than those in CeO2. Then the incorporation of Ca2þ with a larger ionic size, a lower formal charge and a smaller electronegativity than Ce4þ ions as well as F ions with a higher electronegativity than O2 anions into ceria was attempted25 because CaF2 and CeO2 compounds adopt the ˚ 26 and same fluorite network with an almost identical cell parameter equal to 5.46 A 27 ˚ respectively. Furthermore the refractive index of CaF2, around 1.5 in the visible 5.41 A range, and the contribution of fluorine because of its high electronegativity, allow the reduction of the electronic polarizability and consequently the refractive index of ceria. The occurrence of Ca2þ and F ions in the vicinity of Ce4þ and O2 should contribute to modify strongly the absorption and scattering properties at the origin of the evolution of the band gap associated to the 2p (O) ! 4f (Ce) charge transfer involving O2 and Ce4þ ions. In order to reduce the refractive index in the visible range and the electronic polarizability as well as the oxidative catalytic activity associated with the generation of oxygen species, F ions have been substituted for O2 anions as Ca2þ ions for Ce4þ cations. Coprecipitations in highly basic and fluorinated media have been attempted (pH > 12). XRD, Rietveld analysis, chemical analysis, 19F NMR and TEM/EELS measurements have been carried out for an accurate characterization of Ce1xCaxO2x and Ce1xCaxO2x-y/2Fy compounds and to identify the local environments of fluorine.28 The evolution of Ce-X (X ¼ O, F) chemical bonding in this series was described and discussed considering the variation of UV absorption properties and the energy shift of the optical band gap versus

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the chemical composition. However, no photocatalytic and oxidation catalytic activity measurements have been performed on these series. Moreover, in order to overcome the drawbacks of titanium oxide, our approach has also been to increase its ‘anionic electronegativity’, focussing our attention on the modification of the nature and the number of anions in the vicinity of the cations and to decrease the electronic polarizability. In this context, the Ti-O(OH)-F system was investigated.29 Partial substitution of oxygen atoms by fluorine ones is associated with (i) the reduction of the network electronic polarizability and of the dielectric constant; as a consequence, compounds with lower refractive indices should be obtained with a higher transparency in the visible light region and (ii) the occurrence of localized charges inducing the loss of photocatalytic activity. The first-reported titanium oxyfluoride, TiOF2, was synthesized and characterized in 1955 by Vorres and Donohue.30 Focussing on this, in 2006, Reddy et al. 31 investigated the electrochemical behaviour of TiOF2 with Li metal under discharge-charge conditions. Topotactic Li reactions with other ReO3-type compounds have been investigated.32 More recently, Mel’nichenko et al.33 have shown the occurrence of few other titanium oxyfluorides; however, no information about the crystal structure of these compounds was given. In addition other ammonium oxofluorotitanates were prepared and the structures described.34–35 Finally the optical and electrical properties of Ti oxyfluorides adopting the rutile form were published.36 During the synthesis investigation, the modification of the R ¼ [HF]/[Ti] ratio has enabled the preparation of several compounds that adopt different structural types. For increasing R values, compounds deriving from the anatase, the hexagonal tungsten bronze and the ReO3 types have thus been synthesized.37 Among them, the structural properties of the ReO3-derived compound obtained for R ¼ 3, i.e. for the highest fluorine content, are detailed.29 The effect of the substitution of fluoride ions F for O2 ones in the vicinity of Ti4þ cations has been investigated in a first step on the basis of electronic structure calculation, by determining the variation of dielectric function and optical constants, refractive index and attenuation coefficient in UV-visible range.38 These calculations showed that the light-scattering properties of fluorine-substituted phases can be controlled by modifying the fluorine amount and consequently the cell volume. Chemical analysis, TGA and density measurements allow the accurate determination of the chemical formula of Ti hydroxyfluorides. The structural features of the compounds will correlated with their UV-absorption properties.

8.2 8.2.1

Synthesis of Tetravalent Ce and Ti-based Oxyfluorides Preparation of Ce-Ca-based Oxyfluorides

The Ce-Ca based oxyfluorides were prepared by coprecipitation performed in a basic medium at pH > 12. The starting materials are Ce(III) nitrates and Ca chlorides. Yamashita et al.19 have already explored this route but only in the case of Ce-Ca based oxides. In our case, the pH, which is a key parameter, was adjusted taking into account

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the respective domain of precipitation of hydroxides and this route was also applied to oxyfluorides with a fitted stirring duration. The Ce-Ca based oxides synthesis consists of a simultaneous addition in deionized water at around 45 C of Ce(NO3)3 þ CaCl2 aqueous solutions and a 4 M sodium hydroxide solution. After adjusting the pH at values higher than 12, the Ce3þ ions were oxidized into Ce4þ ions using hydrogen peroxyde (30 %). Then the mixture was stirred during 3 h or 20 h. In the case of the preparation of oxyfluorides, after the stirring step, the solution containing Ce-Ca hydroxides was added to a basic fluorinated solution (pH > 12) prepared from HF (40 %) and NaOH solutions with a precursor ratio equal to F/Ca ¼ 2. This mixture was stirred during 1 h, the pH being under control all along the maturation of particles. The precipitate was collected after centrifugation and washed several times with deionized water (pH ¼ 6–7). Then, the precipitates, as recovered, were dried at 100 C during one night and finally annealed under air at 600 C for 12 h. Crystalline compounds with pale yellow coloration were obtained.

8.2.2

Preparation of Ti-based Oxyfluorides

Numerous hydrothermal syntheses have been developed in order to prepare the divided TiO2 phase.39-40 For instance, in our group, several Al-based hydroxyfluorides have been prepared by an original synthesis route, namely the microwave-assisted route41–43 and the structural features have been determined and correlated with the acidic properties. The same method have been adapted to prepare original highly divided Ti(IV)-hydroxyfluorinated compounds.37–38 The peculiar heating mode of microwave irradiation44 leads to a rapid and homogeneous heating that gives rise to numerous advantages as compare with conventional routes. Such advantages can be summarized as follows: increase of the kinetics of reaction, phase selectivity, homogeneous precipitation and fast evaluation of the relevant synthesis parameters. Additionally, because of the interest in such a method, microwave oven technology dedicated to chemistry has been developed over the last decade offering safe equipment and very reproducible experiments owing to the possibility of accurately controlling the temperature, pressure and power. Titanium hydroxyfluorides were synthesized using microwave-assisted processes. Experiments were conducted in a microwave-accelerated reaction system MARS5 (CEM Corporation, Matthews, NC, USA) operating at 2.45 GHz with a power supply varying in the 300–1200 W range, using HP500 vessels. The multimode instrument was equipped with temperature and pressure monitoring device (Pmax 30 bar; Tmax ¼ 200 C). In order to investigate the Ti-O(OH)-F system accurately, several syntheses were performed with a ratio R ¼ [HF]/[Ti] ranging from 0 to 3; TiOCl2.(HCl)x (Ti 15% ; HCl 38 – 42 %) and HF aqueous solution (40 %) were used as precursors. In the case of Ti hydroxyfluorides containing a low Ti3þ content, identified by ESR experiments, isopropanol was preferred to water as solvent, as well as Ti isopropoxide instead of TiOCl2 as precursor and the R ¼ [HF]/[Ti] precursor ratio vary from 3 to 4. The microwave-assisted synthesis could be summarized as a two-step process. The first one, corresponding to an activation period, consisted of three steps: heating up to 90 C at a rate of

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12 C/min, annealing at 90 C for 30 min, and cooling down to room temperature. It should be noted that for synthesis temperatures higher than 90 C, and annealing durations longer than 30 min, TiO2 anatase-type structure was obtained as impurity. At the end of the first step, a limpid solution was obtained. The second step consisted in a microwave-assisted precipitation under drying conditions (with an argon flow and a primary vacuum at 90 C); the resulting white powder was recovered, further washed with alcohol and filtrated using a Millipore Amicon Stirred Ultrafiltration Cell. Finally, the powder was outgassed at 100 C under primary vacuum for four hours.

8.3

8.3.1

Chemical Compositions and Structural Features of Ce and Ti-based Oxyfluorides Elemental Analysis

The Ce/Ca atomic ratio in the Ce-Ca based oxides and oxyfluorides as well as Ce-M-Na based oxides and oxyfluorides with M ¼ Sn or Zr were determined by wavelength dispersive spectrometry (WDS) using CAMECA SX 630 microprobe. The fluorine content was quantified by F titration with a specific electrode. The F/Ti atomic ratio, as well as the homogeneity of the sample, were checked by wavelength dispersive spectrometry (WDS) using a CAMECA SX100 electron probe microanalysis (EPMA) on samples pelletized to get a planar surface. The quantitative determination was performed on the basis of the intensity measurements of the more energetic Ti, O, F X-ray emission lines using TiO2 and TiF3 as reference compounds. Density measurements were performed using either a liquid medium (bromobenzene) or Micromeritics pycnometer operating under helium flow (Micromeritics AccuPyc 1330 Instrument). Electronic Spin Resonance (ESR) spectra were recorded at T ¼ 5 K and at room temperature on an ESR 500 Bru¨cker spectrophotometer working at ¼ 9.449 GHz in order to identify the local environments of Ti3þ. Thermal analyses were conducted in the temperature range 25–800 C using two instruments: (i) a Setaram SETSYS Evolution thermoanalyser and (ii) a simultaneously coupled TGA-MS device, a Netzsch STA 409C Skimmer, equipped with a Balzers QMG 421 and a Pulse TA unit.25, 29, 45

8.3.2

Magnetic Measurements

In order to reach a conclusion about the absence of Ce3þ (4f1) paramagnetic ions, the molar magnetic susceptibility of some Ce-Ca-M (M ¼ Sn or Zr) and Ce-Sn-Na oxides and oxyfluorides compounds were determined from the measurements of the magnetization M (w ¼ M/H) at 300 K with a field of 4 T using the superconducting quantum interference device (SQUID) magnetometer. The magnetization M (M ¼ w*H) was determined on 130–150 mg powdered sample in the temperature range 4–300 K with an applied field (H) of 4 T.

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8.3.3

About the Chemical Composition of Ce1xCaxO2x and Ce1xCaxO2xy/2Fy Series

The composition of Ce-Ca based oxides (Oi samples) and oxyfluorides (Fi samples) are reported in Tables 8.1 and 8.2 respectively, where the Ca/Ce, Na/Ce and F/Ce atomic ratios are given. The Ca/Ce precursor atomic ratios are also listed. In oxides as well as in oxyfluorides, the Ca/Ce atomic ratios determined on the final powder are generally lower than the Ca/Ce precursor ones. This feature illustrates the importance of the pH during the synthesis. Indeed to improve the Ca2þ substitution for Ce4þ ions, the pH must be as high as possible because of the solubility constant of respective hydroxides. The precipitation of Ca(OH)2 begins indeed at pH higher than 11.5. Moreover, it is not excluded that washing with deionized water at pH ¼ 6–7 leads to eliminate Ca2þ ions owing to the high solubility of Ca2þ ions under pH ¼ 7. The F/Ce atomic ratio in Table 8.2 increases with the Ca/Ce atomic ratio in good agreement with the affinity of fluorine for calcium atoms. But whereas the F/Ca precursor atomic ratio is equal to 2, the F/Ca atomic ratio in the final product is always lower than 1. In the oxyfluoride series, two samples F6 and F7 have the same Ca/Ce atomic ratio (0.40) but the F/Ce atomic ratio is higher for F7 (0.52) leading to a mixture of Ca-Ce oxyfluorides and CaF2. In the case of F6, CaF2 was not detected as in the other oxyfluorides with lower Ca contents. This difference illustrates the importance of the duration of the maturation of the oxyhydroxides: 3 h for F7 and 20 h for F6. In the former conditions,

Table 8.1 Sample

Elemental analysis of Ce-Ca based oxides Ca/Ce precursor atomic ratio

Ca/Ce final atomic ratio

Na/Ce final atomic ratio

0 0.04 0.11 0.20 0.28 0.39 0.54

0 0.053 (1) 0.081 (5) 0.207 (4) 0.273 (4) 0.33 (1) 0.414 (4)

0.094 (1) 0.033 (5) 0.018 (4) 0 0 0 0

O1 O2 O3 O4 O5 O6 O7

Table 8.2 Sample F1 F2 F3 F4 F5 F6 F7*

Elemental analysis of Ce-Ca based oxyfluorides Ca/Ce precursor atomic ratio 0.04 0.11 0.18 0.28 0.39 0.54 0.54

Ca/Ce final atomic ratio 0.051 (2) 0.094 (3) 0.156 (6) 0.252 (5) 0.342 (8) 0.410 (4) 0.40 (3)

* mixture containing Ce1xCaxO2xy/2Fy and CaF2

Na/Ce final atomic ratio 0.035 (6) 0.013 (2) 0 0 0 0 0

F/Ce final atomic ratio < 0.01 0.02 (1) 0.03 (1) 0.09 (1) 0.23 (1) 0.34 (1) 0.52 (1)

New Nanostructured Fluorocompounds as UV Absorbers

235

for high Ca content, the Ca2þ ions are probably not sufficiently homogeneously distributed into the network. So, the introduction of fluorine ions in the mixture leads to the segregation of Ca atoms and finally to the formation of CaF2. In the latter conditions, the duration of the oxide’s maturation induced a better homogeneity of Ca2þ ions into the network of ceria. As F ions were introduced into the Ca-Ce matrix, the formation of CaF2 was avoided, leading to Ca-Ce oxyfluorides with high Ca/Ce and F/Ce atomic ratios. Coming from NaOH addition, Naþ ions have been detected in samples with the lowest Ca contents, so that (Ca þ Na)/Ce atomic ratio is always lower than 0.10. It can be thus considered a first step that alkaline and alkaline-earth cations are substituted for Ce4þ cations. In a recent 23Na NMR study, which will be published later, on such Ce-based compounds adopting fluorite-type structure, the recorded spectra present characteristic features of electric-field gradient and isotropic chemical distributions (asymmetry of the central line and broadening of the spinning side bands). This is in agreement with the occurrence of Naþ disordered environments with various numbers of O2 and F anions at different distances surrounding Naþ cations. To confirm the Ce valence state in Ce-Ca based oxides and oxyfluorides after the coprecipitation and the thermal treatment at 600 C under air, magnetic measurements were carried out. In Table 8.3, molar susceptibilities at 300 K of these samples after diamagnetic corrections are listed as well as those of Ce2Sn2O7,46 CeF3, commercial CeO2 from Alfa Products and CeO2 data reported in the literature.47 In CeF3 and Ce2Sn2O7 compounds, cerium ions are totally in trivalent state whereas they are totally in tetravalent state in CeO2. The magnetic susceptibility of cerium trifluoride follows a Curie-Weiss law versus the temperature (100 K £ T £ 300 K) with a theoretical effective magnetic moment equal to 2.54 mB/Ce atom (2.43 mB/Ce atom for the experimental value). CeO2 is assumed to be diamagnetic; its magnetic susceptibility is equal to 0,6.104 emu/ mol and is almost 50 times lower than those of CeF3 and Ce2Sn2O7 compounds. The comparison of their magnetic susceptibilities with paramagnetic and diamagnetic references in Table 8.3 reveals that the Ce3þ content is so low that it cannot be determined. Moreover, the 19F MAS-NMR study, mentioned in the following section also confirms the absence of any dipolar broadening between fluorine ions and Ce3þ ions leading to consider that the majority of Ce ions are in the tetravalent state. Therefore, as the cerium is considered only in tetravalent state in Ce-Ca based oxides and oxyfluorides, the following formula can be proposed for these systems: Ce1xCaxO2x for the oxides and Ce1xCaxO2xy/2Fy for the oxyfluorides. Table 8.3 Room-temperature magnetic susceptibility of Ce-vased oxides and oxyfluorides as well as Ce2Sn2O7, CeF3 and CeO2 compounds as references Compositions O1 Ce0.91Na0.09O1.87 O6 Ce0.75Ca0.25O1.75 F5 Ce0.75Ca0.25O1.67F0.17 Ce2Sn2O7 [24] CeF3 CeO2 [25] CeO2 (Alfa Products)

m (emu/mol) 1,1.104 1,3.104 1,2.104 3,1.103 2,1.103 0,6.104 0,5.104

Functionalized Inorganic Fluorides

236

Thermogravimetry analysis coupled with mass spectrometry TGA-MS was performed on Ce0.75Ca0.25O1.67F0.17 as shown in Figure 8.1. Two main regions were detected. The first weight loss attributed to H2O departure occurs below T ¼ 150 C and the second one is observed around 450 C and 850 C and is attributed to HF and CO2 departures. HF loss starts from 600 C and CO2 from 450 C. The departure of carbon dioxide suggests the presence of carbonate species into the sample associated to Ce4þ and Ca2þ ions as it has already been reported.48 After gas chromatography analysis, the C content was estimated in this sample around 1.6 weight %. Only TGA analyses were performed on Ce0.87Ca0.13O1.86F0.03, Ce0.80Ca0.20O1.77F0.07 and Ce0.71Ca0.29O1.59F0.24 compositions and are presented in Figure 8.2. The two regions DTA/uV

Ion Current *10–10 /A

TG /% 100

[3]

TG

20.0 15.0

95

m18 : H2O

5.0 0

2.0

1.5

↑ exo –5.0

m17 : OH–

80

2.5

10.0

90 DTA 85

3.0

–1.85

–2.04

1.0 –10.0

75

–15.0

m19(x100) : F– 100

200

–20.0

300

400 500 Temperature /°C

600

700

0.5

0

800 DTA/uV

Ion Current *10–10 /A

TG /% 100

[3]

TG –2.04

m44 :

12

CO2+

20.0

1.0

–1.85 15.0 0.8

95

10.0

90 DTA 85

5.0

m45(x50) / 13CO2+

0.6

0

↑ exo

–5.0 80

0.4

–10.0 0.2

75

I(m44) : I(m45) = 90 : 1 12C : 13C = 100 : 1

–15.0 –20.0 0

100

200

300

400

500

600

700

800

Temperature /°C

Figure 8.1 TGA-MS analysis of Ce0.75Ca0.25O1.67F0.17 composition (m ¼ 44, 45: CO2, m ¼ 18, 17: H2O, OH and m ¼ 19:F). Reproduced with the permission of ACS or American Chemical Society

New Nanostructured Fluorocompounds as UV Absorbers

237

0 –1

dm/m (%)

–2 –3 a

–4 –5 –6

H2O departure

b

HF and CO2 departure

–7 c

–8 100

200

300 400 500 Temperature (°C)

600

700

800

Figure 8.2 Thermogravimetry analysis of (a) Ce0.87Ca0.13O1.86F0.03, (b) Ce0.80Ca0.20O1.77F0.07 and (c) Ce0.71Ca0.29O1.59F0.24 compositions. Reproduced with the permission of ACS or American Chemical Society

mentioned above were also observed for these three compounds and it appears that the total weight loss observed at 800 C increases with the Ca content: 3.9, 5.2 and 7.8 weight % for compounds with the molar ratio Ca/Ce ¼ 0.15, 0.25 and 0.41 respectively. Whereas the water loss seems to be always between 2 and 3 weight %, the HF and CO2 rates increase strongly with the Ca content. This feature is in good agreement with the increase of the Ca content in these compositions, associated to the presence of carbonates and fluoride ions.

8.3.4

8.3.4.1

About the Structure and Local Environment of Fluorine in Ce1xCaxO2xy/2Fy Series X-Ray Diffraction Experimental Setup and Refinements

The Ce-based compounds were characterized by X-ray powder diffraction using a Philips PW 1050 diffractometer in a Bragg-Brentano geometry with Cu-K radiation ˚ and K2 ¼ 1.54441 A ˚ ). The intensity data were collected at room tempera(K1 ¼ 1.54059 A ture over a 2 range of 5–110 with 0.02 steps and integration time of 10 s. For some samples, a more accurate study was performed using monochromated Cu-K1 radiation ˚ ). In this case, diffractograms were recorded on a PANalytical X’Pert Pro (K1 ¼ 1.54059 A with Ge(111) incident beam monochromator. Data were collected over a range of 15–120 with 0.017 steps. For each sample, a powder X-ray diffraction analysis was carried out to determine the cell parameter evolution and to control the phase purity. Figure 8.3 shows some diffractograms. The position and intensity of diffraction peaks are in good agreement with a cubic symmetry with the Fm3m space group related to

[440]

[511]

[422]

[331] [420]

[400]

[222]

[311]

[220]

[200]

Functionalized Inorganic Fluorides [111]

238

d

c b

a 100

[331]

[420]

80

[400]

60 2θ (°)

[222]

40

[311]

20

d

c

b

a 55

60

65

70

75

80

2θ (°)

Figure 8.3 Powder X-ray diffraction patterns of (a) Ce0.91Na0.09O1.87, (b) e0.87Ca0.13O1.86F0.03, (c) Ce0.75Ca0.25O1.67F0.17 and (d) Ce0.71Ca0.29O1.59F0.24 compositions. Reproduced with the permission of ACS or American Chemical Society

New Nanostructured Fluorocompounds as UV Absorbers b

239

O2–

c

Ce4+

a

Figure 8.4 Fluorite-type structure of CeO2 with cerium and oxygen environments. Reproduced with the permission of ACS or American Chemical Society

fluorite type structure. This structure can be described (Figure 8.4) as a fcc array of cerium ions in cubic symmetry with oxygens occupying all the tetrahedral sites. The a lattice parameters, refined by the Rietveld method,49–51 are listed in Table 8.4 and ˚ for Ca/Ce ¼ 0 and F/Ce ¼ 0 to 5.4217 (4) A ˚ for Ca/Ce ¼ 0.41 and range from 5.4077 (1) A F/Ce ¼ 0.34. These values are approximately between the cell parameters of CeO2 ˚ )27 and CaF2 (5.46 A ˚ ).26 The evolution of the lattice parameters with the Ca atomic (5.41 A content in oxides and oxyfluorides is illustrated in Figure 8.5 and 8.6, respectively. Both for oxides and oxyfluorides the curves follow a Vegard’s law, thus confirming that Ca2þ have been substituted for Ce4þ ions. The increase of the cell parameters results from the Ca2þ substitution for Ce4þ ions and as a consequence of the larger ionic radius for Ca2þ ˚ and 0.97 A ˚ 52, respectively for eightfold coordination number. compared to Ce4þ : 1.12 A In addition, the higher the Ca content is, the broader the X-ray diffraction lines are patterns as shown in Figure 8.3. Indeed the peaks of the Ce0.71Ca0.29O1.59F0.24 compound are twice broader than those of the Ce0.91Na0.09O1.87 compound. The broadness of

Table 8.4 Lattice parameters a and crystallite size tc determined from powder XRD diffractograms of Ce-Ca based oxides and oxyfluorides i

Sample Oi

˚) a (A

tc (nm)

Sample Fi

˚) a (A

tc (nm)

1 2 3 4 5 6 7

Ce0.91Na0.09O1.87 Ce0.92Ca0.05Na0.03O1.91 Ce0.91Ca0.07Na0.02O1.90 Ce0.83Ca0.17O1.83 Ce0.79Ca0.21O1.79 Ce0.75Ca0.25O1.75 Ce0.71Ca0.29O1.71

5.4077 (1) 5.4096 (1) 5.4122 (2) 5.4144 (2) 5.4171 (2) 5.4183 (3) 5.4204 (3)

15 16 11 10 9 7 7

Ce0.92Ca0.05Na0.03O1.91F0.00 Ce0.91Ca0.08Na0.01O1.90F0.02 Ce0.87Ca0.13O1.86F0.03 Ce0.80Ca0.20O1.77F0.07 Ce0.75Ca0.25O1.67F0.17 Ce0.71Ca0.29O1.59F0.24 Ce1xCaxO2x-y/2Fy þ CaF2

5.4097 (1) 5.4109 (2) 5.4138 (2) 5.4162 (2) 5.4201 (3) 5.4217 (4) 5.4213 (3)

15 9 10 9 9 7 9

240

Functionalized Inorganic Fluorides 5,424 5,422 5,420

a (Å)

5,418 5,416 5,414 5,412 5,410 5,408 5,406 0,00

0,05

0,10

0,15

0,20

0,25

0,30

0,35

0,40

0,45

Ca / Ce atomic ratio

Figure 8.5 Dependence of the lattice parameter on the Ca atomic content in Ce-Ca based oxides. Reproduced with the permission of ACS or American Chemical Society

5,424 5,422 5,420

a (Å)

5,418 5,416 5,414 5,412 5,410 5,408 0,00

0,05

0,10

0,15

0,20

0,25

0,30

0,35

0,40

0,45

Ca / Ce atomic ratio

Figure 8.6 Dependence of the lattice parameter on the Ca atomic content in Ce-Ca based oxyfluorides. Reproduced with the permission of ACS or American Chemical Society

New Nanostructured Fluorocompounds as UV Absorbers

241

diffraction peaks results from two contributions: crystallite size and microstrains. So the crystallite size is very small (< 20 nm) and decreases as the Ca/Ce atomic ratio raises as reported in Table 8.4. The a lattice parameters and tc crystallite sizes were determined by a profile-matching (Le Bail49 fit) using the Thompson-Cox-Hastings function50,54 (function 7 in the Fullprof package) where the contributions for crystallite size and for microstrains are weighted.53 The Thompson-Cox-Hastings function is often used to refine profiles with broad diffraction peaks because it is the more appropriate model for line-broadening analysis where the Lorentzian and Gaussian contributions for crystallite size and for microstrains are weighted.53–54 So in this case, the peak shape is simulated by the pseudo-Voigt function, which is a linear combination of a Gaussian and a Lorentzian function (Table 8.5). ˚ ), (Ce,Ca)-(O,F) distances Table 8.5 Compositions, Ca/Ce atomic ratio, lattice parameters (A ˚ ) and particle sizes tc (nm) of Ce1xCaxO2xy/2Fy samples (A Samples a b ca d e

Formula

Ca/Ce atomic ratio

˚) a (A

Ce0.71Ca0.29O1.59F0.24&0.17 Ce0.75Ca0.25O1.67F0.17&0.16 Ce0.80Ca0.20O1.77F0.07&0.16 Ce0.87Ca0.13O1.86F0.03 0.11 Ce1xCaxO2xy/2Fy&z þ CaF2

0.410 (4) 0.342 (8) 0.252 (5) 0.156 (6) 0.40 (3)

5.4237(2) 5.4233(2) 5.4162(2) 5.4170(2) 5.4244(1) 5.4636(2)

˚) dMX (A 2.3485 2.3484 2.3453 2.3456

tc (nm) 7 9 9 9 8

a For this composition, the lattice parameter and the particle size was determined from diffractogram recorded on a Philips PW 1050 diffractometer in a Bragg-Brentano geometry with Cu-Ka radiations (Ka1 and Ka2).

The refinement of oxide as well as oxyfluoride diffractograms has shown that line broadening is not due to the Gaussian contribution of microstrains (DST*tan ¼ 0). The crystallite sizes vary from 16 to 7 nm and decrease when Ca/Ce atomic ratio raises but are of the same order of magnitude for oxides and oxyfluorides for the same Ca/Ce atomic ratio. The estimated standard deviations corrected with the Berar factor55 are given in parentheses. The stabilization of Ca2þ and F ions into the Ce-based oxides adopting a fluorite-type structure with nanosized particles can then be explained. In many cases, alkaline or alkaline earth as well as fluorine are excellent sintering agents, which contribute to the particle growth. In the Ce1xCaxO2x and Ce1xCaxO2xy/2Fy series, the particle size does not increase and remains around 10 nm. The electronegativity, formal charge and ionic radius are quite different for Ce4þ and Ca2þ ions as those of O2 and F anions. Ca2þ cations are constrained to be sixfold coordinated to oxygen as in CaO56 whereas Ce4þ ions prefer the cubic coordination as in CeO2 leading to the stabilization of anionic vacancies and the formation of defects as well as dislocations. Moreover, the O coordination number varies from 6 in CaO to 4 in CeO2 whereas the F coordination number changes from 4 in CaF2 to 2 in CeF457 and the cation-anion bond distances can vary on a large scale from ˚ to 2.50 A ˚ depending on the nature of ions. Such variations can explain the 2.20 A stabilization of such compositions and the occurrence of strong constrains in the network maintaining nanosized particles. Figure 8.7 shows the evolution of strains with the atomic

242

Functionalized Inorganic Fluorides 7 6

strains (*103)

5 4 3 2 1 0 0,00

0,05

0,10

0,15 0,20 0,25 0,30 Ca / Ce atomic ratio

0,35

0,40

0,45

Figure 8.7 Evolution of microstrains (Dd/d) with Ca/Ce atomic ratio for Ce-Ca based oxides (white circles) and oxyfluorides (black triangles). Reproduced with the permission of ACS or American Chemical Society

ratio Ca/Ce for the Ce-Ca based oxides and oxyfluorides. A recent study by Rodriguez et al.24 has shown that the Ca2þ substitution for Ce4þ ions in Ce1xCaxO2x nanoparticles prepared by microemulsion method induced stress into the lattice due to the differences between Ce4þ and Ca2þ ion features like size, charge and usual coordination number. Then, it has been demonstrated that the strains into the lattice increase with the Ca content and are consequently attributed to oxygen vacancies. This trend is also observed for Ce1xCaxO2x and Ce1xCaxO2xy/2Fy compounds prepared by coprecipitation in basic medium. Indeed Figure 8.7 illustrates that the strains into the structure are more important for compounds with high Ca content. Moreover, the comparison between Ce-Ca based oxides and oxyfluorides with the same Ca/Ce atomic ratio and almost the same crystallite size reveals that the fluorination of the oxides seems to reduce the strains into the network (comparison of samples Ce0.75Ca0.25O1.75 and Ce0.75Ca0.25O1.67F0.17, as well as samples Ce0.71Ca0.29O1.71 and Ce0.71Ca0.29O1.59F0.24). As it has been developed in the previous paragraph, the Ca2þ substitution for Ce4þ leads to the creation of oxygen vacancies into the lattice. These vacancies contribute to the strains formation. The number of vacancies and consequently the strains increase when the Ca/Ce atomic ratio raises and decrease when these anionic vacancies are partially filled with fluorine atoms. 8.3.4.2

CaF2 Observation and Solubility Limit of Ca2þ and F Ions into Ceria

The important width of diffraction peaks and the very close cell parameters of CeO2 and ˚ and 5.46 A ˚ respectively) mean that a XRD analysis with monochromatic CaF2 (5.41 A Cu- K1 radiation is necessary to detect the presence of CaF2 in the sample.

New Nanostructured Fluorocompounds as UV Absorbers

243

b

a b

a 20

40

60

80

100

120

2θ (°)

Figure 8.8 Powder X-ray diffraction pattern of a) Ce0.71Ca0.29O1.59F0.24 and b) Ce1xCaxO2xy/2 Fy þ CaF2 samples obtained with monochromated Cu-Ka radiation (K1 ). Reproduced with the permission of ACS or American Chemical Society

In the same manner, a diffractogram without the Cu- K2 radiation was recorded for Ce0.71Ca0.29O1.59F0.24 composition confirming the absence of calcium fluoride (Figure 8.8). The presence of CaF2 as impurity was also evidenced using 19F MAS NMR. As has been previously mentioned for Ce0.71Ca0.29O1.59F0.24 composition (sample F6), a long maturation of particles of oxyhydroxides (20 h compared to 3 h for the sample F7) prevents the formation of CaF2. A simple ionic model without anionic vacancies shows that the ionic ratio pﬃﬃﬃ rcations =ranions ¼ 3 1 (8:1) corresponds to the theoretical limit of a cubic close packed of cations in which anions occupy tetrahedral holes. The cationic radii in oxygen environment which are present in majority in a ˚ and rCa2þ ¼ 1.12 A ˚ cubic coordination without anionic vacancies are: rCe4þ ¼ 0.97 A (Shannon52) and the O2 radius in fourfold coordination in fluorite-type structure is ˚ (case of CeO2). The expression in Equation (8.1) thus becomes rO2 ¼ 1.38 A x rCe4 þ þ ð1 xÞ rCa2 þ pﬃﬃﬃ ¼ 31 rO2

(8:2)

leading to x ¼ 0.73 for the limit composition. In these conditions, the corresponding composition Ce0.73Ca0.27O1.73 is very close to the Ce0.71Ca0.29O1.71 compound as

244

Functionalized Inorganic Fluorides

˚ which is the common value used prepared. Considering now a O2 ionic radius of 1.40 A for such anions, an x limit value equal to 0.64 is obtained. Then it is difficult to conclude about the Ca limit of solubility. Finally one should have to point out when the Ca/Ce precursor content increases up to 1, a mixture of Ce1xCaxO2x oxide and CaCO3 is obtained. Moreover, in Ce-Ca based oxides, the presence of Ca2þ ions leads to the creation of oxygen vacancies and Ce4þ and Ca2þ cations can adopt various coordination numbers, probably between 6 and 8. On a steric point of view, the occurrence of oxygen vacancies contributes also to reduce the Ca limit of solubility. In the case of fluorinated compounds, the F ionic radius in a tetrahedral coordination is ˚ .52 In CaF2, considering a cubic close always lower than one of O2 ion : rF ¼ 1.31 A ˚ , a rF ¼ 1.25 A ˚ is found. Then the F substitution for O2 packed array with rCa2þ ¼ 1.12 A should increase the Ca solubility limit, thanks on one hand to the theoretical reduction of global anionic radius but also because of the decrease of the number of anionic vacancies. It is then reasonable to consider that the Ca solubility limit is close to 0.3 corresponding both to Ce0.7Ca0.3O1.7 and Ce0.7Ca0.3O1.58F0.25 limit compositions. 8.3.4.3

Evaluation of Cation-anion Bond Lengths in Fluorite-type Network

The a lattice parameters and the tc crystallite sizes of four compositions determined by the Rietveld refinement are listed in Table 8.5. The Ce, Ca and O/F occupancies are fixed to the values found on the basis of chemical analyses. The average (Ce,Ca)-(O,F) bond distances can then be deduced from the refined lattice parameters. One sample, whose elemental analysis has given 29 mol % Ca2þ and 36 mol % F (sample e) appears to be a ˚ and 5.4636(2) A ˚ in good mixture of two phases whose lattice parameters are 5.4244(1) A agreement with the occurrence of Ce-Ca based oxyfluoride Ce1xCaxO2xy/2Fy and CaF2 ˚ )26 respectively. The refinement of this two-phase mixture is illucompound (a ¼ 5.46 A strated in Figure 8.9. The Rietveld refinement does not allow the quantification of the CaF2 content because, as suggested by Brindley et al.,58 the calculated intensity must be corrected by a particle absorption contrast factor. For pure phase samples (Figure 8.9), a good agreement was obtained between the experimental and calculated diffractograms. The Rietveld data resulting from the refinement of X-ray diffractograms corresponding to Ce0.75Ca0.25O1.67F0.17 and Ce0.71Ca0.29O1.59F0.24 compositions are presented in Tables 8.5 and 8.6. The isotropic thermal displacements Biso were refined by fixing the cation and anion occupancies from the elemental chemical composition. In this model, Ce4þ and Ca2þ cations are randomly distributed in the fluorite network (4a Wickoff position) and taking into account the presence of oxygen vacancies, cations can be sixfold, sevenfold or eightfold coordinated to oxygen and fluorine at bond distances ˚ . As the fluorine and calcium contents increase, the isotropic thermal around 2.35 A displacements Biso of both cations and anions seem to slightly increase. This evolution can be explained by the local deviations of cations from Wickoff positions leading to various (Ce,Ca)-(O,F) chemical bondings with different bond lengths, as developed in the next section, and not by the anionic vacancies whose molar concentration remains stable around 0.15 whatever the composition. Ce-F, Ca-F, Ce-O and Ca-O distances in a fluorite type structure were estimated using bond-valence calculation proposed by Brown and Altermatt.59 This model describes the relationship between the

New Nanostructured Fluorocompounds as UV Absorbers

245

bond length (rij) and the bond valence (sij) associated to i and j elements : r r 0 ij Vi ¼ sij ¼ exp (iii) where B is an empirical parameter whose value: j j B is estimated for an oxide as 0.37. Vi represents the oxidation state of the i element and r0 is a constant characteristic of the i-j bonds: here r0 was refined from the reported bond lengths in CeF4,57 CaF2,25 , CeO226 and CaO56 (see Table 8.7). Then, the average rij bond length in the fluorite-type structure can be deduced in order to respect either the cation or the anion valence that can lead to a different result, i.e. a pair of rij values can be proposed for each bond length (Table 8.8).

a b 10

30

50

70

90

110

130

2θ (°)

a b 10

30

50

70

90

110

130

2θ (°)

Figure 8.9 X-ray diffraction patterns of Ce0.75Ca0.25O1.67F0.17 (top), Ce0.71Ca0.29O1.59F0.24 (middle) and Ce1xCaxO2xy/2Fy þ CaF2 (bottom) samples: (a) observed (...), calculated (—) signals and (b) difference diagram; the tick marks represent the Bragg position of the diffraction lines. Reproduced with the permission of ACS or American Chemical Society

Functionalized Inorganic Fluorides

246

a

b 10

30

50

70

90

110

130

2θ (°)

Figure 8.9 (Continued)

˚ 2), occupations and reliability Table 8.6 Atomic positions, isotropic thermal displacements (A factors of Ce0.75Ca0.25O1.67F0.17 and Ce0.71Ca0.29O1.59F0.24 samples. Estimated standard deviations multiplied by Berar’s factor are indicated in parentheses cRp ¼ 8.42

Ce0.75Ca0.25O1.67F0.17

cRwp ¼ 7.88

RBragg ¼ 1.82

Atoms

Site

x

y

z

˚ 2) Biso(A

Occupancy (%)

Ce Ca O/F

4a 4a 8c

0 0 0.5

0 0 0.5

0 0 0.5

0.44(2) 0.44(2) 1.16(8)

0.75 0.25 0.92

cRwp ¼ 6.67

RBragg ¼ 1.64

cRp ¼ 6.87

Ce0.71Ca0.29O1.59F0.24 Atoms

Site

x

y

z

˚ 2) Biso(A

Occupancy (%)

Ce Ca O/F

4a 4a 8c

0 0 0.5

0 0 0.5

0 0 0.5

0.61(2) 0.61(2) 1.35(9)

0.71 0.29 0.92

˚ ), cationic/anionic coordination number and Brown Table 8.7 Bond lengths (A ˚ ) estimation in CaO, CeO2, CeF4 and CaF2 binary and Altermatt r0 constant (A systems Compound

Chemical bond

Bond lengths ˚) (A

CeO227 CeF457 CaF226 CaO56

Ce-O Ce-F Ca-F Ca-O

2.343 2.200–2.304 2.365 2.404

[Cation] /[ anion] coordination numbers 8/4 8/2 8/4 6/6

˚) r0 (A 2.0870 1.9866 1.8541 1.9965

New Nanostructured Fluorocompounds as UV Absorbers

247

˚ ) in Table 8.8 Estimated M(Ce/Ca)X(O/F) chemical bond lengths (A fluorite network based on the Brown-Altermatt model59 M-X

˚) Bond length (A (to respect the cation valence)

˚) Bond length (A (to respect the anion valence)

Ce-O Ca-O Ce-F Ca-F

2.34 2.51 2.24 2.37

2.34 2.25 2.50 2.37

Nevertheless, for systems with equilibrated effective charges, i.e. for Ce-O and Ca-F distances in CeO2 or CaF2 fluorite-type structures, only one rij value is obtained per compound. Otherwise, in a hypothetical fluorite-type structure with nonequilibrated charges as in ‘CeF2’ or ‘CaO2’ systems, Ce4þ cations require Ce-F bond lengths of ˚ ) and in the same way, ˚ whereas F anions need higher Ce-F distances (2.50 A 2.24 A 2þ ˚ Ca ions in eightfold coordination can be predicted at 2.51 A from oxygen ions whereas ˚ is calculated for O2 ions in Ca2þ tetrahedral coordination. a Ca-O distance at 2.25 A Thus, several opposite trends have to be taken into account when CeO2 is partially substituted by CaF2 and it appears difficult to predict the average Ce-F and Ca-O bond lengths inside Ce1xCaxO2xy/2Fy fluorite type structure. Nevertheless, in all cases, the occurrence in the same compound of Ce-F and Ca-O bonds will induce some disequilibrium either on the cation side, or on the anion side, so these chemical bonds will induce some constrains and relaxations. On the contrary, because CeO2 and CaF2 are both ˚ and 2.365 A ˚, equilibrate charge systems with roughly the same bond lengths (2.343 A respectively), one can already predict a segregation of the fluoride ions around the calcium cations; indeed it is the best way to limit bond discrepancies inside Ce1xCaxO2xy/2Fy compounds. ˚ and in Finally, one should notice that the Ce-F average bond length value is about 2.5 A ˚ , i.e. the bond lengths predicted from the a same way, Ca-O bond lengths at about 2.25 A anionic centre. Indeed, because the tetrahedral coordination sphere of the anions is less flexible than the eight-fold coordination sphere of the cations, it was considered that the respect of the anions valence governs the bond lengths. 8.3.4.4 19

19

F NMR Analysis and Local Environment of Fluorine in Fluorite-type Network

The F MAS NMR spectra were recorded on an Avance 300 Bruker spectrometer (7 T) with a Larmor frequency of 282.2 MHz for 19F, using a high speed CP MAS probe with a 2.5 mm rotor. The external reference chosen for the isotropic chemical shift determination was C6F6 (isoC6F6 versus CFCl3 ¼ 164.2 ppm60). The spectra were acquired using a single pulse sequence (1 ms), followed by the free-induction decay acquisition. The delay between two acquisitions was 1 s. For such a delay, the quantitativity of the spectra was checked. The discrimination of isotropic peaks from spinning sidebands was achieved by recording spectra at two different spinning frequencies (20 and 25 kHz). The reconstructions of the 19F NMR spectra were performed with the DMFIT software61. In this study, the iso values, the relative intensities and the line widths are the relevant parameters.

248

Functionalized Inorganic Fluorides 4

3

2

1

CaF2

CeF4

(a)

(b)

(c)

(d)

300

200 100 δiso (ppm)

0

Figure 8.10 19F MAS NMR spectra of (a) Ce0.71Ca0.29O1.59F0.24, (b) Ce0.75Ca0.25O1.67F0.17, (c) Ce0.80Ca0.20O1.77F0.07 and (d) Ce0.87Ca0.13O1.83F0.03 samples at 25 kHz. The number of scans is equal to 2k for samples (a) and (b), 8k for sample (c) and 36k for sample (d). The twofold arrow points out the 19F diso value range in CeF4. The sticks indicate the four isotropic 19F NMR lines. The width of the sticks is proportional to the diso variation with composition. Reproduced with the permission of ACS or American Chemical Society

With the aim of studying the local environments of fluorine ions and more particularly by identifying the distributions of Ca2þ or Ce4þ cations around fluorine ions located in 8c sites, 19F NMR measurements were undertaken on the Ce1xCaxO2xy/2Fy samples. The 19F MAS NMR spectra recorded at 25 kHz are reported in Figure 8.10. The signalto-noise ratio is low for all the samples and decreases from spectrum a to spectrum d, in agreement with the related decreasing low F contents. The reconstructions of the spectra, recorded at 20 and 25 kHz, were achieved with four or three, for the lowest Ca content, broad lines between 60 and 190 ppm. An example is given in Figure 8.11. The isotropic chemical shift values, relative intensities and line widths deduced from the reconstruction are presented in Table 8.9. The line widths do not depend on the spinning frequency, which indicates that heteronuclear and homonuclear dipolar interactions are fully averaged at 20 kHz. Then, the line widths of the NMR resonances, which increase with the Ca, F and vacancy contents mirror the isotropic chemical shift distributions. This may be related to the occupancy of Ce4þ and Ca2þ cations as well as O2 and F anions on the same crystallographic sites (4a and 8c respectively) and the occurrence of

New Nanostructured Fluorocompounds as UV Absorbers

249

3

2

1

4

300

200

100

0

δiso (ppm)

Figure 8.11 Experimental (top) and calculated (bottom) 19F MAS NMR spectrum of Ce0.71Ca0.29O1.59F0.24 at 20 kHz. The unnumbered lines are the spinning side bands. Reproduced with the permission of ACS or American Chemical Society

Content of fluorine ion environment (%)

7 6 5

FCaCe3 FCa2Ce2 FCa3Ce FCa4

4 3 2 1 0 0.00

0.02

0.04

0.06 0.08 Fluorine content

0.10

0.12

0.14

Figure 8.12 Content of the different types of fluorine environments as a function of F atomic ratio in the Ce1xCaxO2xy/2Fy samples. Reproduced with the permission of ACS or American Chemical Society

anion vacancies, which lead to various F-Ce and F-Ca bond lengths. The relative intensities of the NMR lines depend on the Ca and F contents as shown in Figure 8.12. The distinct 19F resonances (at 63–65 ppm, 113–116 ppm, 150–160 ppm and 180–190 ppm) were assigned to different types of environments, FCa4, FCa3Ce, FCa2Ce2, and FCaCe3,

250

Functionalized Inorganic Fluorides

using the superposition model, proposed by Bureau et al.60 This model takes into account the cation-fluorine bond distances and assigns isotropic chemical shifts to fluorine crystallographic sites through the use of phenomenological parameters. The proportions of fluoride ion environments determined by the 19F NMR study (Table 8.9) are clearly different from those predicted by the statistical calculations on the basis of a random distribution of anions in the 8c sites and cations in the 4a sites (Table 8.10). The contents of the fluoride ion environments are equal to the relative line y intensities (Table 8.9) multiplied by 2 (y corresponds to the F molar content). The comparison of these values shows that the more numerous are the Ca2þ cations in the vicinity of F ions, the higher is the discrepancy between experimental and predicted contents. Based on the bond lengths determined from the Brown and Altermatt model (Table 8.8), the average value of the (Ce,Ca)-(O,F) distances in the FCa4nCen environments can be calculated with a linear combination of F-Ca and F-Ce bond lengths. ˚ when n increases from 0 to 4. Moreover, These distances increase from 2.37 to 2.50 A this latter distance appears underestimated considering isotropic chemical shift calculations (see above). The more numerous the Ce4þ cations in the vicinity of F ions, the larger the average F-(Ce,Ca) distances, the stronger the constrains for the network and the less stable the environments of fluoride ions. This is confirmed by the absence of FCe4 sites that would create too many constrains for the network of the studied Ce1xCaxO2xy/2Fy compounds where the average distance (Ce,Ca)-(O,F) is ˚. around 2.35 A The observed contents of the fluoride ion environments may be understood as the result of a compromise between the strong affinity of fluoride ions in fourfold coordination for Ca2þ cations in the fluorite structure, the F and Ca contents and the network constrains induced by longer Ce-F bonds. High F and Ca contents lead to the segregation of the F and Ca2þ ions into the network of the studied Ce-Ca oxyfluorides and finally to the formation of CaF2 as observed in sample e which has the same Ca/Ce ratio as sample a (Table 8.5) but has been prepared as samples b, c and d (see experimental section). Since the proportions of FCa4 environments are similar for samples a and b, it may be inferred that a longer duration of maturation of the oxyhydroxides (sample a), allows the synthesis of a sample with a more homogeneous distribution of Ca2þ cations corresponding to a larger content of FCa4 avoiding thus the formation of CaF2 during the fluorination. The highest F content of sample a can also be explained by the increase of the proportions of the more stable environments for fluorine ions (FCa3Ce and FCa2Ce2) and the decrease of the proportions of the less stable environment for fluorine ions (FCaCe3) compared with sample b. Table 8.9 shows that for samples a and b corresponding to the highest Ca and F contents, the proportion of FCa4 environments remains stable around 13 %. Moreover, the proportions of FCa2Ce2 and FCa3Ce sites tend towards 50 % and 25 % of the total F sites respectively. Then, the proportion of FCa4 sites seems to reach a limit rate from the atomic ratio Ca/Ce ¼ 1/3 corresponding to the Ce0.75Ca0.25O1.67F0.17 composition whereas the anionic vacancy content reaches a limit equal to 1/12 starting from Ca/Ce atomic ratio around 1/4 with a lower F content (Table 8.5). Then, the Ca and F contents as well as the proportion of each fluorine environment seem to be independent of the creation of anionic vacancies into the ceria matrix.

Table 8.9 Isotropic chemical shift values (ppm), relative intensities (%) and line widths (ppm) of the 19F NMR lines, with assignments, as deduced from NMR spectrum simulations for Ce1xCaxO2xy/2Fy samples Samples

a b c d Attribution

line 1

line 2

line 3

line 4

iso (–1)

I (–0.5)

width (–0.5)

iso (–1)

I (–0.5)

width (–0.5)

iso (–1)

I (–0.5)

width (–0.5)

iso (–1)

I (–0.5)

width (–0.5)

63.5 64.0 65.0 /

12.9 13.0 2.8 0.0

11.0 9.0 6.0 /

113.5 114.5 116.0 115.5

28.3 24.4 24.8 25.8

22.0 21.0 18.0 13.5

150.5 152.5 156.0 160.5

47.8 45.2 67.6 71.1

24.0 24.5 25.0 22.0

179.0 181.5 185.0 188.0

11.0 17.5 4.9 3.1

16.0 18.0 10.0 6.5

FCa4

FCa3Ce

FCa2Ce2

FCaCe3

252

Functionalized Inorganic Fluorides

Table 8.10 Predicted and experimental (italics) contents (%) of fluoride ion environments for the studied Ce1xCaxO2xy/2Fy samples (the total F content is equal to y/2 considering that there are twice anionic sites than cationic sites) Samples a b c d

8.3.4.5

FCa3Ce

FCa4 0.085 0.033 6.103 0.00

1.55 1.10 0.10 0.00

0.83 0.40 0.090 0.011

3.40 2.07 0.87 0.39

FCa2Ce2 3.05 1.79 0.54 0.12

5.74 3.84 2.37 1.07

FCaCe3 4.98 3.59 1.43 0.51

1.32 1.49 0.17 0.046

FCe4 3.05 2.69 1.43 0.86

0.00 0.00 0.00 0.00

TEM and EELS Analysis: Particle Size and Composition of Nanoparticles

Transmission electron microscopy (TEM) pictures were recorded using a microscope TECNAI F20 equipped with a field emission gun, operating at 200 kV with a point resolution of 0.24 nm. The powder in toluene suspension was placed on a copper mesh and dried in air. The pictures were realized in bright field. Moreover, in order to explore accurately the composition homogeneity of each grain, electron energy-loss spectroscopy (EELS) spectra were recorded with a Philips CM20 microscope operating at 200 kV with a point resolution of 0.27 nm. The analyses were performed from the calcium L2,3-edge, fluorine and oxygen K-edges and cerium M4,5-edge after subtracting the background. The samples were crushed in an agate mortar and dispersed on a copper grid covered by a holey-carbon-film. The particle size and the morphology were studied by TEM on Ce0.75Ca0.25O1.67F0.17 (sample F5). The morphology is not homogeneous whereas the particle sizes are quite identical. Figure 8.13 shows particles of about 10 nm, which is similar to the crystallite size determined by XRD analysis. The images lead to the conclusion that the particles are monocrystalline. Figure 8.14 illustrates EELS spectra realized on two isolated particles, showing that the four elements have been detected even if the F-K edge is very low because the F/O atomic ratio is close to 0.1. The Ca/Ce average atomic ratio was evaluated on isolated nanoparticles to 0.30 with a standard deviation of 0.04. This value corresponds to the ratio deduced from microprobe analysis: Ca/Ce ¼ 0.342(8). Furthermore, EELS analysis on isolated nanoparticles of Ce0.75Ca0.25O1.75 oxide reveals also the quite good homogeneity of the oxide with a molar ratio Ca/Ce ¼ 0.29 and a standard deviation of 0.04. EELS experiments have confirmed the good Ca/Ce homogeneity at the nanoparticle level on both oxide and oxyfluoride systems.

8.3.5 8.3.5.1

Composition and Structure of Ti-based Hydroxyfluoride Microwave Synthesis

As already mentioned in Chapter 2 of this book, the [HF] content is a relevant parameter since it allows the stabilization of several phases in an aqueous medium. Focusing on the R ¼ [HF]/[Ti] ¼ 3 molar ratio, the synthesis operated at T ¼ 90 C with Ti oxychloride as precursor enables the preparation of a single phase related to TiOF2 type structure as revealed by X-ray diffraction analysis with crystallite size around 80 nm. Interestingly, an

New Nanostructured Fluorocompounds as UV Absorbers

253

Figure 8.13 High-resolution TEM images (200 kV) of Ce0.75Ca0.25O1.67F0.17 composition, agglomerated particles (top) and single particle (bottom). Reproduced with the permission of ACS or American Chemical Society

increase in the reaction temperature and/or the duration time leads to the occurrence of TiO2 anatase as an impurity while an excess of fluoride ions is present. A lower synthesis temperature, i.e. 90 C, thus favours the coexistence of F and OH/O2 species in the vicinity of Ti4þ, whereas additional energy favours the formation of the oxide through oxolation reaction. The effect of the nature of the synthesis medium has also been considered through the replacement of both titanium precursor and solvent. An organic medium has therefore been used containing titanium isopropoxide as precursor and isopropanol as solvent. In these conditions, two different R ¼ [HF]/[Ti] molar ratios have been studied, i.e. 3 and 4. The corresponding X-ray diffraction powder patterns (not shown) are similar to those obtained from the aqueous medium with a broadening of the X-ray peaks suggesting smaller particle size. Here, the salient point is that the recovered powders are coloured, which is consistent with mixed-valency Ti. Such a point will be discussed later. 8.3.5.1.1 X-Ray, Neutron and Electron Diffraction Structural properties of titanium hydroxyfluorides were characterized by X-ray and neutron powder diffraction. For X-ray characterization, diffraction data were recorded on a

254

Functionalized Inorganic Fluorides

Ca L2–3 Ce M4–5

Intensity (a.u.)

OK FK

300

400

500

600 700 800 Energy Loss (eV)

900

1000

1100

Figure 8.14 EELS spectra showing the Ce M4,5, Ca L2,3, O and F K-edges of Ce0.75Ca0.25O1.67F0.17 composition recorded on two single grains. Reproduced with the permission of ACS or American Chemical Society

PANalytical X’Pert PRO diffractometer in the Bragg-Brentano geometry using a Ge (111) ˚ ) and equipped with incident beam monochromator (Cu-K1 radiation; l ¼ 1.54056 A X’Celerator detector. For neutron characterization, data collection were realized at ILL ˚ ). In order to minimize (Grenoble) on the CRG-D1B line using the wavelength (l ¼ 1.287 A inelastic scattering of hydrogen atoms, isotope exchange was performed using D2O as solvent. The diffraction data were analysed using the Rietveld technique as implemented in the Fullprof Suite Program. Peak shape was described by a pseudo-Voigt function, and the background level was fitted with linear interpolation between a set of given points. Transmission electron microscopy (TEM) experiments were performed with a JEOL 2000 FX microscope operating at 200 kV. 8.3.5.2

Composition and Structural Features

In the X-ray diffraction powder pattern of the R ¼ 3 compound (in aqueous medium) most of main lines fall at similar 2 values as those of TiOF2 published by Vorres and Donohue.30 These authors successfully synthesized this compound using two different routes: hydrolysis of titanium tetrafluoride or titanium trifluorochloride or reaction of aqueous or anhydrous HF with TiO2. The oxyfluoride crystallizes in the Pm–3m space ˚ , and a powder density of 3.06 g.cm3 with one group with a lattice parameter of 3.80 A TiOF2 molecule per unit cell. The crystal structure is built up of corner-sharing Ti(X)6 (X ¼ O, F) octahedra, where the O and F atoms are randomly distributed. Therefore, in a first attempt, the structure of R ¼ 3 compound was refined using a similar model, corresponding to a full occupancy of each site and a TiOF2 formula. However, such hypothesis led to significant differences between calculated and observed XRD peak intensities, especially for the (100) and (110) reflections (Figure 8.15). This mismatch

New Nanostructured Fluorocompounds as UV Absorbers 100

Intensity (a.u.)

(a)

Observed Calculated Difference Bragg positions

200 210

311 222 320

110 111

20

40

211 220 300 310

60 80 2-Theta λCuKα1 (°)

100

Intensity (a.u.)

(b)

400 410 321 411

100

120

Observed Calculated Difference Bragg positions

200 210

311 222 320

110 111

20

255

40

211 220 300 310

60 80 2-Theta λCuKα1 (°)

400 410 321 411

100

120

Figure 8.15 X-ray powder diffraction analysis (Rietveld) of Ti hydroxyfluoride considering two hypotheses in the Pm–3m space group: (a) without Ti vacancies, (b) with Ti vacancies. Reproduced with the permission of ACS or American Chemical Society

could be correlated with a high value of isotropic thermal displacement parameter for the ˚ 2) ˚ 2) and a negative value for the X one (X ¼ O, F ; B ¼ –0.14(4) A Ti site (B ¼ 2.17(3) A (Table 8.11). The strong dependence between site occupancy and thermal displacement parameters led us to improve the structural model by adding titanium vacancies and fixing B factors to acceptable values. Finally both these parameters were refined simultaneously. The best fit

256

Functionalized Inorganic Fluorides

Table 8.11 Atomic positions, isotropic thermal displacements, occupancies and interatomic distances determined by powder XRD (T ¼ 293K) data analysis for Ti hydroxyfluoride (deuter˚) ated phase) with Pm 3m space group (a ¼ 3.81017(4) A Atom

Site

Empirical Formula TiOF2 Ti O/F

1a 3d

x Rp ¼ 0.212

y

z

Rwp ¼ 0.201

0 0

˚ 2) B (A

ai Rbragg ¼ 0.100

0 0

0 ½

1 1

2.17(3) 0.14(4)

Distances: ˚ Ti – O/F 1.9051(2) A ˚ O/F – O/F 2.6942(2)A

Table 8.12 Atomic positions, isotropic thermal displacements, occupancies and interatomic distances determined by powder XRD (T ¼ 293 K) data analysis for Ti0.75 0.25(OH,D)1.5F1.5 ˚ ) taking into account Ti (deuterated phase) with Pm-3m space group (a ¼ 3.81012(4) A vacancies Atom

Site

x

y

Rp ¼ 0.181

Empirical Formula Ti0.75 0.25(OH,D)1.5F1.5 Ti O/F

1a 3d

0 0

z

0 0

0 ½

˚ 2) B (A

ai Rwp ¼ 0.164

Rbragg ¼ 0.0372

0.74(1) 1

1.06(2) 2.55(6)

Distances: ˚ Ti – O/F 1.9051(2) A ˚ O/F – O/F 2.6942(2)A

is achieved for a refined titanium occupancy value of 0.74(1) and thermal displacement parameters of 1.06(2) and 2.55(6), for titanium and oxygen/fluorine respectively. The observed and calculated diffraction profiles using the ‘titanium vacancies model’ are drawn in Figure 8.15. The R factor and atomic positions are gathered in Table 8.12; ˚ and 2.6942(2) for Ti-O/F and O/F-O/F, refined bond distances are equal to 1.9051(2) A respectively. These Ti-X distances are in good agreement with those generally encountered for the Ti atom. A representation of the structure is given in Figure 8.18. The unit cell is built of corner-shared Ti-X6 octahedra. Experimental evidence for such a structure containing titanium vacancies can be correlated with an experimental density value smaller than the one of pure TiOF2 (i.e. 3.06 g.cm3). Indeed, using two different density measurement methods: bromobenzene pycnometry and helium pycnometry, similar values were obtained: 2.628(3) and 2.631(2), respectively. In order to take into account 25 % of titanium vacancies, a formula close to Ti0.75(OH)1.5F1.5 should be considered. This formula was confirmed by Microprobe analysis measurements showing the presence of 39.6 % of titanium atoms and 30.6 % of fluorine atoms, which are in good agreement with the 39.9 % and 31.6 % calculated values. However, in order to reach conclusions about such chemical formula, TEM investigations have been carried out (Figure 8.17). The accurate observation of the electronic

New Nanostructured Fluorocompounds as UV Absorbers

257

diffraction patterns along [100] [21-1] and [11-1] zone axes (fourfold and sixfold axes characteristic of cubic symmetry) shows the presence of extra spots in between the most intense ones, suggesting the existence of a superstructure with a doubled a-parameter ˚ ). Assuming this cell, all visible spots can be indexed and two reflection (a ¼ 7.60 A conditions were deduced: 0 k l: k þ l ¼ 2n and h 0 0 : h ¼ 2n. Therefore, only one space group belonging to m-3 Laue class, Pn-3m can be considered. Regarding this last space group (Pn-3m), four atomic positions were found: 2a (0,0,0) and 6d (0, ½,½) for Ti4þ, 12d (¼,0, ½) and 24k (x, x, z) with x 0.027 and z 0.246 for O2/F. The 24k position is half-occupied instead of a full occupancy for all others positions. Both regular and distorted octahedra are observed too. One should note that trials with 12g (0, 0, x) atomic positions instead of 24k (x, x, z) ones did not bring any improvement. A random distribution of fluorine atoms in 24k positions (50 % occupied) corresponds to a better representation of the structure of Ti0.75(OH)1.5F1.5. The Ti vacancy rate deduced from the Ti occupancy in 2a and 6d sites is equal to 0.15 instead of 0.25 as deduced from chemical analysis. Such a difference remains high but taking into account the number of structural parameters, one should have to consider that the Ti occupancy spread into two sites cannot be determined accurately. Furthermore, the Debye-Waller ˚ 2 for O/F atoms. ˚ 2 for Ti and 1.38/1.22 A factors are correct: around 0.9 A Refined parameters as well as atomic positions and displacements are presented in Table 8.13. Experimental and calculated X-ray and neutron diffraction patterns with this Table 8.13 Atomic positions, isotropic thermal displacement, occupancies and interatomic distances determined by neutron diffraction (T ¼ 293 K) data analysis for ˚ ) taking Ti0.75&0.25(OH,D)1.5F1.5 (deuterated phase) with Pn-3m space group (a ¼ 7.6177(2) A into account Ti vacancies. Atom

Site

x

Ti0.75 0.25(OH,D)1.5F1.5 (Neutron) Ti1 Ti2 O1/F1 O2/F2

2a 6d 12f 24k

0 0 ¼ 0.027(1)

Distances: ˚ [x6] Ti1 – O2/F2 1.9044(3) A ˚ [x4] Ti2 – O1/F1 1.9095 A ˚ [x2] Ti2 – O2/F2 1.9503(3) A

y Rp ¼ 0.015 0 ½ 0 0.027(1)

z Rwp¼0.021

ai

˚ 2) B (A

Rbragg ¼ 0.0278

0 ½ ½ 0.246(3)

0.87(3) 0.80(3) 1 0.50

0.9(2) 0.9(3) 1.38(3) 1.22(4)

˚ O1/F1 – O1/F1 2.7004 A ˚ O1/F1 – O2/F2 2.59(2) – 2.88(2) A ˚ O2/F2 – O2/F2 2.37(3) – 2.95(3) A

new hypothesis are compared in Figure 8.16. The relationships between ReO3-type ˚ ) and the supercell (SG: Pn-3m, a ¼ 7.61 A ˚ ) are represtructure (SG: Pm-3m, a ¼ 3.80 A sented in Figure 8.18. The supercell exhibits two Ti and O/F atomic positions characterizing the ordering of both Ti4þ cations and OH/F anions, in the as-prepared Ti hydroxyfluoride. The structure contains two regular octahedra Ti1(2a) ˚ ) and six distorted octahedra Ti2(6d) (d(Ti2 – O1/F1) ¼ 1.904 A ˚ (d(Ti1 – O2/F2) ¼ 1.905 A ˚ [x2]). Considering these last octahedra, a pseudo [x4] and d(Ti2 – O2/F2) ¼ 1.951 A sequence of [Ti2(O1/F1)4(O2/F2)2] elongated octahedron with [Ti1(O2/F2)6] regular octahedron can be easily visualized.

258

Functionalized Inorganic Fluorides 16 000 002 13 000

Yobs Ycalc Yobs–Ycalc Bragg_position

Intensity (a.u.)

10 000 7000

4000

004 204 202

1000

224

222

404 424

–2000 –5000 2θ (°) 160 000 222

140 000

Yobs Ycalc Yobs–Ycalc Bragg_position

120 000

Intensity (a.u.)

100 000 202 80 000

226 004

60 000

224 40 000

002

404

206

204

444

426

20 000 0 –20 000 10.5

20.5

30.5

40.5

50.5

60.5

70.5

80.5

90.5

2θ (°)

Figure 8.16 X-ray (top) and neutron (bottom) data refinements of Ti hydroxyfluoride in the ˚ ) taking into account vacancies in Ti sites. Reproduced with Pn–3m space group (a ¼ 7.6177 A the permission of ACS or American Chemical Society

Based on the ‘Ti0.75(OH)1.5F1.5’ formula deduced from chemical analysis and this polyhedron sequence, two extreme formulations can be proposed: [TiF6/2]2[Ti2/3&1/3(OH)4/2F2/2]6 [Ti(OH)6/2]2[Ti2/3&1/3(OH)1/3F8/3]6 In order to probe the anionic environment of Ti atoms, the valence bond model has been used. Based on the Ti-F and Ti-O average bond distances in TiF4 and TiO2 rutile-type ˚ , r0(O) ¼ 1.815 A ˚ ) can be calculated. The first structure, two r0 values (r0(F) ¼ 1.723 A hypothesis with [TiF6/2]2[Ti2/3&1/3(OH)4/2F2/2]6 distribution should be considered. In this case the valence of the first Ti atom surrounded by six fluorine atoms in F2 positions is equal to þ 3.67 whereas the second Ti atom surrounded by four hydroxyls in OH1 atomic positions and two fluorine atoms in F2 positions is equal to þ 4.23. Then the deviation from the þ 4 theoretical value for both valence bonds leads to the consideration

New Nanostructured Fluorocompounds as UV Absorbers

259

1 – 11 0 – 11 011 011

[100]

[21-1]

101 011

[11 - 1]

100 nm

˚) Figure 8.17 TEM study of Ti0.75&0.25(OH,D)1.5F1.5 phase considering a supercell (a ¼ 7.6177 A in Pn3m space group with reflection conditions 0kl: k þ l ¼ 2n and h00: h ¼ 2n. A TEM micrograph of Ti hydroxyfluoride showing the size particle is given. Reproduced with the permission of ACS or American Chemical Society

of the occurrence of OH groups in the F2 position as well as F ions in the OH1 atomic position. As far as the valence bonds of F1/OH1 and F2/OH2 atomic positions are concerned, taking into account Ti vacancies in Ti2 atomic position, F atoms in F2/OH2 position as well as OH groups in OH1/F1 atomic position, the F2 valence is equal to – 0.97 and the valence of OH1 is equal to 1.04, both close to the theoretical value 1. Then, the Ti vacancies seems to be exclusively distributed on Ti2 sites despite the results of XRD and neutron diffraction data refinement showing the distribution of Ti vacancies on both Ti1 and Ti2 sites. Furthermore, one should have to consider the occurrence of the distribution of F/OH in two anionic sites. Thus, the electronegativity difference between the two anions OH and F induces a local distortion of Ti4þ cations, creating an ordering. In terms of the driving force of this new atomic arrangement, the conflict between the small ionic radius of fluoride ions and its lower ligand field leads Ti4þ ions to be surrounded by various distributions of OH/F anions. Off-centred Ti4þ ions in regular octahedra will be surrounded mainly by F ions, even if a OH/F distribution can occur on these sites, whereas elongated octahedra contain OH/F mixed anions, the two ligand fields creating anisotropy with Ti-vacancies stabilized in these distorted octahedral sites. Although it is rather unusual to observe a segregation between fluoride and oxide octahedra in oxide fluorides; this property is found rather often in hydrated fluorides in which MF6 octahedra alternate with [M(OH2)]6 octahedra, as in AlF3.3H2O, in chain compounds RbMnF4.H2O and CuSiF6.4H2O, or in inverse weberite Fe2F5.2H2O. Concerning the cationic vacancies, one should note that the K2NiF6 structure63 can also be described as

260

Functionalized Inorganic Fluorides

b c

a

Ti

b

F/O

c

Ti a

F/O

Ti1

b c

a

Ti2

Ti1 Ti2 OH/F F/OH

˚ ) and Figure 8.18 Relationships between ReO3-type structure (SG: Pm3m, a ¼ 3.8088 A ˚ & Ti0.75 0.25(OH,D)1.5F1.5 with the supercell (SG: Pn3m, a ¼ 7.6177A); The two Ti sites and the two O/F atomic positions characterizing the occurrence of Ti4þ cations and O2/F anions ordering are pointed out. Reproduced with the permission of ACS or American Chemical Society

deriving from the perovskite (elpasolite) or ReO3-type networks by an ordering between cationic vacancies and NiIV ions. It has been pointed out that new compositions containing small amounts of Ti3þ can be prepared using microwave-assisted synthesis and a large HF/Ti precursor ratio (3
Thermal Behaviour

The thermal behaviour of Ti0.75(OH)1.5F1.5 was evaluated by simultaneously coupled thermogravimetric and mass spectrometry analysis (TGA-MS).

New Nanostructured Fluorocompounds as UV Absorbers

261

100

80

R (%)

60

40

20 d-d intervalencies bands 0 300

400

500 600 Wavelength (nm)

700

800

Figure 8.19 UV-Visible spectra of Ti hydroxyfluorides prepared by microwave assisted synthesis using either water, R ¼ HF/Ti ¼ 3, Ti oxychloride as precursor (red curve), or isopropanol as solvent, Ti isopropoxide as precursor, R ¼ 3 (green curve), or isopropanol as solvent, Ti isopropoxide as precursor, R ¼ 4 (blue curve). Reproduced with the permission of ACS or American Chemical Society

1 106 Room Temperature * 600

8 105

5K

5

Intensity (a.u.)

6 10

4 105

g = 1.90 (Ti3+)

5

2 10

Site D4h (g⊥, g ⁄⁄ ) D4h site (T = 5 K)

0 –2 105 –4 105 –6 105 2500

3000

3500

4000

4500

Magnetic Field (G)

Figure 8.20 ESR spectra (T ¼ 5K and room temperature) of Ti hydroxyfluoride prepared by microwave assisted synthesis using isopropanol as solvent, Ti isopropoxide as precursor and R ¼ HF/Ti ¼ 3 precursor ratio. Reproduced with the permission of ACS or American Chemical Society

262

Functionalized Inorganic Fluorides 0

100

–1

90

8 10–9

–2

13 wt.% 80 TG (wt.%)

1 10–8

–3

6 10–9

TG (wt.%) 70 DTG (wt.%/min)

45 wt.%

–4

–5

60 m18 m19

2 10–9 –6

50

–7

100

200

300 400 500 Temperature (°C)

600

700

800

0

Ion Current /A

0

DTG (wt.%/min)

40

4 10–9

Figure 8.21 Thermogravimetry analysis coupled with mass spectrometry of Ti0.75OH1.5F1.5 compound showing the water and TiF4 evolutions. Reproduced with the permission of ACS or American Chemical Society

The thermogravimetric curve displayed in Figure 8.21, shows two distinct weight losses. These two separate weight losses are associated on the differential thermogravimetric curve with a broad peak extending from 100 to 320 C, and a sharp one centred at 570 C. The absence of weight loss below 140 C excludes the existence of free water (H2O). Analyses of the evolved gases by mass spectrometry indicates that the first and second weight losses are mainly related to the departure of water vapour and of fluorine, respectively. Besides, in order to understand the huge second weight loss (around 45 %), a departure of Ti- based fluoro-species coupled with that of fluorine should be considered. Unfortunately the m/z signals relating to TiF4 species cannot be detected because most of the gas is condensed at the entrance of mass spectrometer. Based on the results of thermogravimetric analysis coupled with mass spectrometry, one can summarize the scheme of Ti0.75(OH)1.5F1.5 decomposition as follows:

Temperature domain ( C)

Released species

Resulting products

25–140 140–350

Stable H2O (13 wt. %)

400–600

F and TiF4 (45 wt. %)

Ti0.75(OH)1.5F1.5 Ti hydroxyfluorides or ‘TiOF2’ þ " TiO2 anatase TiO2 anatase

New Nanostructured Fluorocompounds as UV Absorbers

263

The observed weight losses are in agreement with the theoretical ones. Indeed, in a first step, a complete loss of OH groups (as 0.75 H2O) corresponds to a theoretical weight loss of 15 wt. %. In a second step, the decomposition of Ti-based fluoro-species (as 0.375 TiF4) is associated with a theoretical weight loss of 52 wt. %, slightly higher than the observed 45 wt. % weight loss. The small discrepancies reported previously could be explained by a more complex decomposition scheme, assuming a simultaneous departure of OH groups and fluorine as HF. This decomposition scheme is also in agreement with the ex situ annealing of Ti0.75(OH)1.5F1.5 during 4 hours under N2 atmosphere. The evolution of the X-ray powder pattern is reported in Figure 8.22. After the first weight loss, the XRD pattern of 101

Ti0.75(OH)1.5F1.5ReO3 Thermal decomposition - 350°C Thermal decomposition - 600°C

Intensity (a.u.)

200

* Al sample-holder ** WLα

004

h' k' l'- TiO2 Anatase h k l - Ti0.75(OH)1.5F1.5

105 211 204

103

112 213

**

220 116

215 301 224 008

* 110

200 210 111

30

40

211

220

300 310

222

100 20

50

60

70

80

90

2-Tetha λCuKα (°)

Figure 8.22 XRD patterns of ReO3-derived Ti hydroxyfluoride annealed under N2 during 4 h at various temperatures (T ¼ 350 C, T ¼ 600 C). Reproduced with the permission of ACS or American Chemical Society

the sample annealed at 350 C exhibits mainly Bragg peaks corresponding to TiOF2 (ICDD 08-0060)27 Bragg peaks corresponding to TiO2 anatase can also be noted (ICDD 21-1272)64. Finally, after the second weight loss, XRD pattern of sample annealed at 600 C shows only the presence of Bragg peaks indexed with the TiO2 anatase structure.

8.4

UV Shielding Properties of Divided Oxyfluorides

Diffuse reflectance measurements were carried out in the UV-visible range from 200 to 800 cm1 on a Cary spectrophotometer at room temperature in order to illustrate the UV-shielding properties of such compounds.

264

8.4.1

Functionalized Inorganic Fluorides

The Ce-Ca-based Oxyfluorides Series and UV-shielding Properties

The diffuse reflectance spectra of various oxides: Ce0.83Ca0.17O1.83, Ce0.79Ca0.21O1.79 and Ce0.75Ca0.25O1.75, and oxyfluorides: Ce0.87Ca0.13O1.86F0.03, Ce0.80Ca0.20O1.77F0.07 and Ce0.75Ca0.25O1.67F0.17 are given in Figure 8.23. One has to point out the interesting UV-shielding properties of these compounds with an absorption edge around 400 nm and the lower refractive index and high transparency in the visible range in the case of Ce1xCaxO2xy/2Fy compositions taking into account the Gladstone-Dale relationship with n(CeO2) ¼ 2.50, n(CaO) ¼ 1.80 and n(CaF2) ¼ 1.50 at l ¼ 550 nm.

100 90 diffuse reflection (%)

80 70 60 50

Ce0.83Ca0.17O1.83

40

Ce0.79Ca0.21O1.79

30

Ce0.75Ca0.25O1.75

20 10 0 300

350

400

450 500 550 wavelength (nm)

600

650

700

100 90 diffuse reflection (%)

80 70 60 50

Ce0.87Ca0.13O1.86F0.03

40

Ce0.80Ca0.20O1.77F0.07

30

Ce0.75Ca0.25O1.67F0.17

20 10 0 300

350

400

450 500 550 wavelength (nm)

600

650

700

Figure 8.23 UV-visible diffuse reflectance spectra of Ce-Ca based (top) oxides and (bottom) oxyfluorides. Reproduced with the permission of ACS or American Chemical Society

New Nanostructured Fluorocompounds as UV Absorbers

265

In the oxide series, the absorption edge relative to the charge transfer band (CTB, 2p (O) ! 4f (Ce)) shifts to lowest energies as the Ca content increases. On the other hand, in the oxyfluoride series, the absorption edge moves to the highest energies as the Ca content raises. However in a previous paragraph, it has been demonstrated that the fluorine content in the oxyfluoride series follows the Ca amount and the higher the fluorine content, the higher the optical band gap associated with the 2p (O) ! 4f (Ce) CTB. The comparison of the optical band gap in CeO2-fluorite and in SrCeO3-perovskite is equal to 3.1 eV (400 nm) and 3.5 eV (360 nm) respectively13 clearly shows that this evolution is governed by the Ce-O chemical bonding, i.e. the coordination number and the Ce-O bond distances. In CeO2-fluorite, the Ce4þ ions are eightfold coordinated to ˚ , whereas in perovskite SrCeO3, Ce4þ ions are sixfold oxygen atoms and Ce-O ¼ 2.34 A ˚ , with an increase in the CTB energy. coordinated to oxygen atoms and Ce-O ¼ 2.25 A Such a trend is associated with the destabilization of 4f (Ce) band and with the strong stabilization of 2p (O) band as both coordination number and Ce-O bond distances decrease. In the oxide series, the progressive introduction of Ca2þ cations into the network contributes to increase the oxygen vacancy rate, and thus reduces the average cationic coordination number. On the other hand, the higher the Ca content, the higher the a-cell parameter relating to the expansion of the framework and the average bond distances. Then it is reasonable to consider that in the oxide series, Ca2þ cations tend to occupy sixfold ˚ leaving Ce4þ coordinated sites as in CaO where Ca-O bond distances lie around 2.40 A 2þ cations with higher coordination number than Ca cations and large Ce-O bond distances. In other terms, in Ce1xCaxO2x oxides, Ca2þ cations force Ce4þ cations to become less electropositive than in CeO2 by adopting a coordination number smaller than 8 and larger bond distances. Thus in the oxide series, the higher the Ca content, the less electropositive the Ce4þ cations, leading to the reduction of the energy of 2p (O) ! 4f (Ce) CTB. In the oxyfluoride series, considering the same amount of Ca2þ ions, the number of anionic vacancies is smaller because of the presence of fluorine ions substituting for O2 ions and the band gap energy varies in the opposite way to the trend observed in the oxides series. For Ce1xCaxO2xy/2Fy compositions, Ca2þ ions because of their strong preference for F ions, occupy probably eightfold coordinated sites with a large amount of fluorine surrounding Ca2þ cations, leaving Ce4þ ions with a lower coordination number and fluorine in its vicinity. Then Ca2þ and F ions constrain Ce4þ cations to be more electropositive than in CeO2, leading to stabilize more the 2p(O) valence band with regard to the 4f(Ce) level. The higher the Ca and F contents in this series, the more pronounced the ionic character of Ce-O chemical bonding, leading to an increasing of the optical band gap. It has thus been demonstrated how a high electropositive cation such as Ca2þ and a high electronegative anion such as F are able to tune the 2p(O) ! 4f(Ce) CTB energy or the band gap energy in Ce-Ca based oxides. F ions being mainly bounded to one (FCa3Ce1) and two (FCa2Ce2) Ce4þ cations whatever the chemical composition, a large number of Ce sites are affected by fluorine. The occurrence of F ions in the vicinity of Ce4þ cations and the increase of the proportion of Ce atoms bound to fluorine show that the Ce-(O,F) chemical bonding becomes increasingly ionic as Ca2þ and F ions content raises. In other words, the global electropositive character of Ce4þ cations increases with Ca and F contents leading to the progressive shift of the optical band gap to higher energy associated to the 2p (O) ! 4f (Ce) charge transfer. Moreover, because of the occurrence of Ce-F and

266

Functionalized Inorganic Fluorides

Ca-F ionic bonds, the electronic polarizability should be reduced then contributing to the attenuation of scattering in visible range. 8.4.2

Ti Hydroxyfluoride and UV-shielding Properties

In order to evaluate the UV-shielding properties of Ti0.75(OH)1.5F1.5 with the one of anatase- and rutile-type forms of TiO2, the respective diffuse reflectance are compared in Figure 8.24. These spectra exhibit an absorption edge at a wavelength in the 380-400 nm range corresponding to the charge transfer O(2p)-Ti(3d). UV Area

Visible Area

100

Reflectance (%)

80

60

40 TiO2 Anatase 20

TiO2 Rutile Ti0.75(OH)1.5F1.5 ReO3

0 250

300

350

400 450 Wavelength (nm)

500

550

Figure 8.24 Comparison of the UV-Visible reflectance of TiO2 and Ti hydroxyfluoride. Reproduced with the permission of ACS or American Chemical Society

In the visible part at 500 nm, the lower reflective intensity for Ti0.75(OH)1.5F1.5 with ReO3-type structure (around 85 %) could be related to the texture and morphology of the powder that exhibits larger size particles (around 80 nm) with respect to the 50 nm size for nano-anatase and nano-rutile. In order to gain a better understanding of the differences between these compounds, the UV-Visible absorption properties have been discussed on the basis of a schematic band diagram. The electronic band structures of titanium oxides and titanium hydroxyfluorides exhibit similar features. The conduction band is characterized by the 3d states of titanium and the valence band is composed of the 2p orbitals of O and/or F atoms. In all these compounds, titanium is surrounded by six oxygen or fluorine atoms in an octahedral environment. The band gap is mainly dependent on three parameters: • Firstly, Ti-O bond distances play a key role; it is well known that longer Ti-O bond distances lead to smaller band gap energy associated with the oxygen-metal transfer and

New Nanostructured Fluorocompounds as UV Absorbers

267

this trend has been quoted in the literature as the ‘red-shift phenomenon’ in the semiconductor band gap. In perovskite-type compounds Sr(Ba)TiO3, which adopts the same ˚ octahedra framework as ReO3-type structure, a Ti-O bond distance around 2.00 A 65 (BaTiO3) is related to a charge transfer band around 3.5 eV . Then, a higher band gap energy should be normally expected in the case of Ti0.75OH1.5F1.5 compound where ˚. the Ti-O is around 1.90 A • In these compounds, three types of octahedra frameworks are found: edge-sharing, corner-sharing, and simultaneously edge and corner-sharing; the presence of edgesharing leads to strong Ti(3d)-Ti(3d) interactions, with short Ti-Ti bond distances of ˚ and 2.96 A ˚ for anatase and rutile forms, respectively. By taking into account 3.05 A competitive bonds, this effect may induce a larger energetic O(2p)-Ti(3d) charge transfer in the anatase form. This phenomenon explains the higher band gap of 3.20 eV for anatase-type structure TiO2, as compared to 3.1 eV for the rutile-type one. • The third parameter corresponds to the presence of mixed anions like, for example, F, O2, OH. The presence of fluoride ions with a higher electronegativity than that of oxygen, may lead to a decrease of the network polarizability and a reduction of the refractive index and finally a more energetic band gap is expected. Taking into account the above points, the occurrence of a charge transfer band around 3.2 eV in the case of Ti hydroxyfluoride could be explained by the presence of OH groups, the reduction of electronic density around O atoms being due to H atoms and the nonbonding character of 2p O orbitals at the top of the valence band. However, the occurrence of Ti vacancies because of the reduction of O-Ti interactions leads to a decrease of bandwith of the valence band as well as the stabilization of O(2p) valence band. Then only the stabilization of OH groups in the vicinity of Ti4þ cations can explain the evolution of absorption edge in the case of Ti hydroxyfluoride.

8.5

Conclusion

New Ce1xCaxO2xy/2Fy compositions have been prepared from coprecipitation of CeCa salts in a basic fluorinated medium (pH > 12). Magnetic measurements lead to conclusions about the absence of Ce3þ paramagnetic species whereas TGA measurements can illustrate the affinity of these compounds with water and carbonates, which increase together with Ca content. The experimental solubility limit corresponds to 30 % of Ca2þ ions both for oxides and oxyfluorides, close to the theoretical solubility based on ionic model with a cubic close packed array of cations with anions in tetrahedral holes. In the case of the oxyfluoride series, the high stability of CaF2 is the limiting factor for the incorporation of Ca2þ and F ions into the matrix. Size particles are around 10 nm and accurate EELS measurements demonstrate the rather good homogeneity in chemical composition of nano-sized particles. The presence of Ca2þ ions in the CeO2 fluorite network allows the creation of oxygen vacancies whose rate can be reduced by the presence of F ions substituting O2 ions with a lower charge and a higher electronegativity. Moreover the Ca2þ substitution in both these oxide and oxyfluoride series leads to an expansion of the unit cell and of the average

268

Functionalized Inorganic Fluorides

Ce(Ca)-O(F) bond distances, despite the formation of oxygen vacancies, with respect to the larger ionic size of Ca2þ ions in eightfold coordination. Furthermore, the F amount increases with increasing Ca contents, illustrating the strong affinity of fluorine for Ca atoms. In this series, F atoms are tetrahedrally coordinated with Ce and Ca atoms located at vertices. Based on the Brown-Altermatt model, the F anions force Ce4þ cations to be ˚ . Then ˚ whereas O2 anions constrain Ca2þ cations to be linked at 2.25 A bounded at 2.50 A ˚ the average bond length is around 2.35 A and corresponds to the Ce-O and Ca-F distances in CeO2 and CaF2 respectively. 19F MAS NMR spectroscopy was used to study the local structure and fluoride ion environments. Four distinct 19F resonances have been observed. They were assigned, using the superposition model proposed by Bureau et al.60 to four different types of environment for the fluorine ion: FCa4, FCa3Ce, FCa2Ce2 and FCaCe3. The contents of these environments are clearly different from those calculated on the basis of a random distribution of anions in the 8c sites and cations in the 4a sites. It shows that, in these Ce-Ca oxyfluorides adopting the fluorite-type structure, F anions have a great affinity for Ca2þ cations leading to an increase of the F amount with the Ca content. The absence of FCe4 environment is explained from steric argument or from steric restriction: the F-Ce bond lengths in the network are too short to accept four fluorine ions in the vicinity of Ce4þ cations. A critical composition was identified corresponding to Ca/Ce ¼ 1/3, a situation where the probability to get FCa4 environment becomes high. Finally, the UV absorption properties show a reduction of the optical band gap as Ca content increases in the Ce1xCaxO2x series confirming results previously found.16 The optical band gap increases with increasing Ca and F contents in the Ce1xCaxO2xy/2Fy series. The noteworthy evolution of the O (2p) ! Ce (4f) charge transfer band illustrates the versatility of Ce4þ ions to form various chemical bonding. In the case of oxides, the presence of Ca2þ ions forces Ce4þ ions to become less electropositive than in CeO2 with an increase of the Ce-O bond lengths and a reduction of the band gap. In the case of oxyfluorides, the presence of both Ca2þ and F ions constrains Ce4þ ions to become more electropositive than in CeO2 with an increase in the optical band gap. Considering the various proportions of FCe4nCan environments and the chemical composition, one can point out that the Ce-(O, F) chemical bonding exhibits a more ionic character as the Ca and F contents increase. Consequently, the absorption threshold shifts to UV range. Then the development of ionic bonds into ceria contributes therefore to the reduction of both electronic polarizability and refractive index. Thus new inorganic UV filters with nanosized particles exhibiting a white colour and broad absorption edge at the frontier between UV and visible range as well as a reduced refractive index in visible range can be thus designed. A new Ti hydroxyfluoride has been prepared by microwave-assisted synthesis. All performed characterizations: chemical analysis, density and TGA measurements, FTIR and powder XRD allow conclusions to be drawn about the occurrence of Ti vacancies in a ReO3-type structure as well as fluorine substituting hydroxyl groups. All experimental data account for the chemical formula Ti0.75&0.25OH1.5F1.5. In a second step, TEM studies reveal the occurrence of superstructure of ReO3 network. Refinements of the powder XRD and neutron diffraction patterns in the Pn-3m space group support the presence of isotropic TiX6/2 octahedra and elongated TiY4/2X2/2 (X,Y ¼ OH/F) octahedra. Based on valence bond calculations, the Ti vacancies seem mainly located in the latter distorted site, whereas

New Nanostructured Fluorocompounds as UV Absorbers

269

F/OH are randomly distributed on the two anionic sites. Moreover Ti3þ can be stabilized in this network by changing the synthesis conditions: the precursor Ti isopropoxide instead of [Ti] oxychloride, the R ¼ [HF]/Ti precursor ratio and the solvent, isopropanol instead of water. ESR measurements confirm the occurrence of elongated octahedra taking into account the stabilization of mixed F/OH anions around Ti3þcations. Finally Ti0.75OH1.5F1.5 exhibits interesting UV-shielding properties with an absorption edge around 3.2 eV as TiO2 anatase form but with a lower refractive index (n ¼ 1.9 in visible range). The optical band gap remains low despite the low Ti-O/F bond distance and can be explained by considering the stabilization of OH groups in the vicinity of Ti4þ cations, which allows the destabilization of the 2p (O) valence band with nonbonding character because of the presence of protons and Ti vacancies. It can be added that, using the same microwave-assisted route with lower R ¼ F/Ti ratios, two other Ti-based frameworks 37 [hexagonal tungsten bronze and anatase forms] have been synthesized. A new generation of Ti-based UV absorbers with low refractive index due to the presence of fluorine has been thus prepared and characterized. Moreover one should note that the titanium hydroxyfluoride Ti0.75OH1.5F1.5 does not exhibit any photocatalytic activity because of the stabilization of F ions in the vicinity of Ti4þ cations and the limitation of charge carriers in such ionic compounds, which is an important feature for applications in UV protection.

Acknowledgement We gratefully acknowledge the European Community for the financial support in the STREP FUNFLUOS (FUNctionalized FLUOrideS) network (NMP3-CT-2004-505575).

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9 Oxyfluoride Transparent Glass Ceramics Michel Mortiera and Ge´raldine Dantelleb a

Laboratoire de Chimie de la Matie`re Condense´e de Paris, UMR CNRS 7574 – Chimie ParisTech, 11, rue Pierre et Marie Curie, 75231 Paris cedex 05, France b Laboratoire de Physique de la Matie`re Condense´e, UMR CNRS 7643 – Ecole Polytechnique, Route de Saclay, 91128 Palaiseau Cedex, France

9.1

Introduction

Oxyfluoride glass-ceramic systems are ambivalent materials that, although mainly oxide glasses, can offer optical properties of fluoride single-crystals when they are doped with rareearth ions. They are often called nanocomposite materials. Their peculiar character is obtained by a classical melting and quenching preparation in air followed by an adapted thermal treatment due to the ability of fluoride phases to crystallize at temperatures lying between the glass transition and the crystallization of the oxide glass matrix. Consequently, the choice of the oxide glass composition and the fluoride phase composition is the key factor to succeed in obtaining glass-ceramic materials. Thermal analysis offers an irreplaceable characterization tool for the glass system and provides all the characteristic temperatures and parameters that govern the stability and devitrification of the glass. Transparent glass-ceramics are obtained only with nanometer-size particles because of the step index between the two separated phases. To get crystal-like optical properties, optically active ions have to be entirely segregated inside the crystal particles. Such a partition is not automatic during nucleation and growth processes and is often partial and somewhat fortuitous.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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In this chapter, we describe different transparent oxyfluoride systems doped with lanthanide ions. We examine the conditions that can lead to a perfectly controlled crystallization process and also the conditions that insure an efficient segregation of the rare-earth ions inside crystallites. Finally, we discuss the optical properties of glass-ceramics, in terms of absorption, emission and lifetimes. We establish a comparison between their optical properties and those of glasses and single-crystals.

9.2

Synthesis

There are three main methods for the synthesis of oxyfluoride glass-ceramics: • The partial devitrification of an oxyfluoride glass is the oldest and best known method. A glass is first made by melting polycrystalline oxide and fluoride precursors. The glass obtained is then annealed at a temperature enabling its partial crystallization. This temperature is usually determined by doing the differential thermal analysis of the glass [1, 2]. The whole synthesis can be performed in air. • The annealing of an amorphous matrix obtained by soft chemistry can also lead to the formation of glass-ceramics. This method has been developed with success by S. Fujihara et al. [3]. In the particular example of the SiO2:LaF3 system, a gel is obtained by reacting the different precursors (tetramethylorthosilicate, La(CH3COO)3, trifluoroacetic acid (TFA)) in an aqueous medium. A Si–O–Si network is formed by hydrolysis. The gel is dried around 100 C to complete the hydrolysis and then it is heated in the temperature range 300–800 C. Above 300 C, lanthanide salts decompose to form LaF3. The resulting material is a glass-ceramic made of a silica gel encapsulating LaF3 crystallites with a size varying between 10 and 30 nm. • Finally, the third method consists of first preparing nanocrystalline fluoride particles and then dispersing them in an amorphous oxide medium. Fluoride particles can be synthesized by various techniques (colloidal synthesis, micellar route and so forth). The last two methods, based on a soft chemistry approach, have two main advantages over the glass devitrification: • First, the temperature required for a sol-gel synthesis is much lower than for a glass synthesis: typically, a glass synthesis requires temperatures around 1200 C for germanate glasses, 1600 C for silicates, while the synthesis temperature doesn’t exceed 800 C using one of these soft chemistry methods. • Second, it allows various shapes, as a gel is much easier to shape than a glass. However, as the soft chemistry route requires the use of organic solvents and inorganic precursors, the as-made materials contain organic groups, which can be very harmful for optical applications. In particular, the presence of OH or CH groups should be avoided, as their vibrations induce nonradiative de-excitations, leading to a reduction of the luminescence of the material [4]. In general, in order to make glass-ceramics for optical applications, the glass devitrification method will be chosen. This process, being the most common, will be described in detail below.

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9.2.1

275

Synthesis by Glass Devitrification

The synthesis of glass-ceramics by devitrification is a two-step method. The first step of the process consists of synthesizing a glass by melting crystalline powders and fast quenching them. The glass obtained is then partially devitrified, i.e. some of its components crystallize whereas the rest remains amorphous. The resulting composite material is a glass-ceramic. The devitrification process can be provoked by a thermal treatment [1, 5], by irradiation with a femtosecond laser [6] or by the combined use of a UV laser and thermal annealing [7]. Although a thermal treatment is mostly used, the use of a laser is spreading as it has the advantage of devitrifying the glass locally, allowing the construction of waveguides in a bulk glass. Two devitrification processes are possible: the nucleation/growth process and the spinodal decomposition, leading to glass-ceramics with very different morphologies. To understand the difference between those two processes, an analogy can be made with the demixion of two liquid phases [8] and one can examine a liquid-liquid phase diagram of a binary system A-B (Figure 9.1). At high temperature, the system A-B forms a homogeneous liquid whatever the composition of the system. While cooling down the system, the liquid A and the liquid B are not miscible any more (area limited by the solid line on Figure 9.1): two phases are distinct. If the system is in the grey area, it is metastable. The phase separation occurs by the nucleation/growth process. On the other hand, if the system is in the dashed area, it is in an instable state and the two phases appear by spinodal decomposition. Those two devitrification processes are described hereafter.

Temperature

Homogeneous liquid Metastable system

Instable system

0 (A)

Composition xB (% mol)

100 (B)

Define the area where the two phases are not miscible

Figure 9.1 Phase diagram of a binary system A-B. The solid line represents a region where Phase A and Phase B are not miscible

Spinodal decomposition is the least encountered process. To minimize its energy, the glass modifies its composition slightly and progressively, until two distinct phases appear (Figure 9.2). After the phase separation, one of the phases spontaneously crystallizes. Usually, the crystalline phase represents more than 70 % of the total volume of the material and has a dendritic shape, as shown in Figure 9.3. This process was for example encountered in some fluorozirconate glasses, with composition of 68.6ZrF4:18.4LaF3: 5AlF3:5GaF3:3ErF3 (mol %) [9].

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(a) Spinodal decomposition C2

C0 distance

distance

distance

t1

t2 > t1

t3 > t2

C1

distance

distance

distance

t1

t2 > t1

t3 > t2

Figure 9.2 Representation of the composition of the crystallites as the devitrification occurs (a) by spinodal decomposition and (b) by a nucleation/growth process. C0 corresponds to the initial composition of the glass, C1 corresponds to the composition of the glassy phase after devitrification and C2 to the composition of the crystallites (b) Nucleation/growth process

(a) Spinodal decomposition

Glass Crystallites

Figure 9.3 Schematic morphology of a glass-ceramic formed (a) by spinodal decomposition and (b) by a nucleation/growth process

The nucleation/growth process, most commonly encountered, consists of two separate steps. The first one is the nucleation, corresponding to the formation of seeds, also called nuclei, whose chemical composition is different from that of the initial glass, thus forming a new phase. The composition of this new phase remains constant over the whole nucleation/growth process (Figure 9.2). The nucleation can be homogeneous in the glass or nonhomogeneous if it occurs around some impurities or from the surface. A seed, which can be defined by its radius r, evolves with time. It can either grow if its size is over the critical size, or disappear otherwise. Indeed, in the case of a homogeneous nucleation, the seed free energy DGseed can be described by two terms: • a negative term ( 34 pr 3 DGcrist , where DGcryst is negative and represents the crystallization free energy) corresponding to the crystallization (i.e. ordering) and inducing a decrease of the system free energy. • a positive term (4pr 2 , where is the surface energy at the interface, per unit area), corresponding to the surface creation (i.e. expansion of the particle). DGseed ¼

4 3 pr DGcrist þ 4pr 2 3

(9:1)

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Free energy ΔG

4πr2 σ

ΔG*

4πr3ΔG 3

r*

Seed radius, r

Figure 9.4 Evolution of the free energy of a spherical seed, as a function of its radius. In dashed lines are represented the volume and surface contributions

The evolution of the seed free energy DGseed as a function of the seed radius r is presented in Figure 9.4. For small values of r, DGseed is governed by the surface energy and thus is positive, resulting in the seed instability. Consequently, the seed disappears. For high r values, the volume term is dominant, leading to a thermodynamically stable seed. Hence, to sustain, the seeds must have a radius bigger than a critical value r*. To reach this critical value, it is necessary to provide some energy DG* to the system. The second step of the process is the growth, where the stable seeds grow at the expense of the small ones. The mechanism is also known as the Ostwald ripening [10]. The growth of the particles, also called crystallites, depends on the diffusion speed of the atoms in the glass and on the interface glass/crystallites [11]. To some extent, it is possible to control the formation of the particles by controlling the initial glass composition, as well as the annealing temperature and time. The resulting morphology of glass-ceramics is illustrated in Figure 9.3. It consists of crystalline particles of a spherical or ovoid shape, dispersed more or less homogeneously in the glassy matrix.

9.2.2

Transparency

The control of the size of the particles in glass-ceramics is fundamental as it plays a key role in the transparency of the glass-ceramics. Indeed, according to Rayleigh’s theory, the intensity of light scattered by a particle follows the relation: 2 r 6 M 2 1 (9:2) I I0 4 2 l M þ2 where I is the scattered light, I0 the incident light, r the radius of the particle and M is the ratio of the refractive index of the particle and the glass. Hence, to get a transparent

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material, at least one of the two following statements should be true: r6 l4

1, i.e. the particle size should be much smaller than the incident light wavelength.

M 1, i.e. the difference between the refractive index of the particles and the one of the glass should be weak. In general, the diameter of the particles must be smaller than a tenth of the wavelength l, i.e. 2r< 10l 40nm, if the refractive index difference between both media (glass and particles) is smaller than 0.3. As the particle size is of great importance, it needs to be carefully monitored during the devitrification process, using X-Ray diffraction (XRD) for instance. Take the example of the GeO2:PbF2:PbOþErF3 system where devitrification is induced by thermal treatment. The temperature and the time of annealing are key parameters for the size of the particles. In Figure 9.5.a., one can observe the evolution of the XRD diagrams of 50GeO2:10PbF2:40PbOþ2ErF3 glass-ceramics after a 10-hour annealing at various temperatures. The intensity of the diffraction peaks, corresponding to the crystallization of the cubic -PbF2 phase and superimposed on the diffuse scattering of the amorphous phase, increases with an increase in the annealing temperature. It indicates the progressive crystallization of the fluoride phase. If the annealing temperature is too high or the annealing time is too long, not only the fluoride phase crystallizes but also some additional oxide phases appear as shown in the case of the SiO2:Al2O3:SrF2þYbF3þPrF3 system in Figure 9.5b. (b)

(a)

SrF2 SrAI2SiO8 Si (Standard)

690°C (200)

Intensity / arb.unit

(111) (311)

(220) GC 380°C GC 370°C GC 360°C GC 355°C

680°C

670°C 660°C (111) 650°C (200)

(220)

(311) (222)

GC 250°C 600°C

Glass

As–made

20

25

30

35

40 45 2θ (°)

50

55

60

65 20

50 30 40 Diffraction angle 2θ /°

60

Figure 9.5 (a) Evolution of the X-ray diffraction diagrams of a GeO2:PbF2:PbOþErF3 glass and the corresponding glass-ceramics annealing for 10 hours at various temperatures. The crystalline phase is b-PbF2. Figure reproduced from reference 12. Copyright (2006) Elsevier Ltd (b) Evolution of the XRD diagrams of a SiO2:Al2O3:SrF2þYbF3þPrF3 glass and glass-ceramics obtained after 8-hour annealing at various temperatures. Figure reproduced from reference 13. Copyright (2007) Ceramic Society of Japan

Oxyfluoride Transparent Glass Ceramics

9.3

279

Different Systems

Glass-ceramics can be classified and differentiated according to the nature of the glassy phase, the crystalline phase and the doping ions. Among the most common glassy phases, one can find silica, silica/alumina or germanate. Changing the glassy phase from silica to germanate has an influence on the conditions of the initial glass synthesis. A silica glass has a melting temperature around 1600 C, whereas a germanate glass can be made at a much lower temperature (1200 C). It also modifies the refractive index of the amorphous phase (from n ¼ 1.401.55 for silica [14] to n ¼ 1.6 –1.7 for germanate [15]). Regarding the crystalline fluoride phase, fluorite-type compounds MF2 (M ¼ Ca, Cd, Cd/Pb, Pb, Sr, Ba), as well as LaF3, are commonly used to make oxyfluoride glassceramics. The fluorite structure, whose one unit cell is represented in Figure 9.6, consists of a cubic network comprising cations M2þ with a coordinance of 8 and twice as many anions F, with a coordinance of 4. The incorporation of a trivalent doping ion, e.g. Ln3þ, requires the substitution of one of the M2þ cations. However, the system is left with an excess of positive charge. To compensate, several mechanisms have been proposed, such as the creation of cation vacancies, following the equation 3M2þ 2Ln3þ þ Pb vacancy [16] or the presence of an interstitial fluoride anion F [17, 18, 19]. Depending on the position of the interstitial fluorine ion, the Ln3þ symmetry changes: • if the interstitial F is situated in another unit cell than the one containing Ln3þ, the Ln3þ local symmetry can be considered as cubic; • if the interstitial F is in the same unit cell as Ln3þ, it induces a distortion from the cubic symmetry, leading for example to a tetragonal (C4v) symmetry if the interstitial F is situated in the <100> direction or to a trigonal (C3v) symmetry if it is situated in the <111> direction, as represented in Figure 9.6.b. All trivalent rare-earth ions have been introduced in glass-ceramics with the aim of developing them as optical materials. Nd3þ, Pr3þ, Er3þ ions are among the most commonly studied. The divalent europium ion has also been introduced. Glass-ceramics are often doped with several rare-earth fluorides. Indeed, codoping glass-ceramics with several lanthanide ions have proved to be very efficient to improve their optical properties, as it will be explained in the section on multi-doped glass-ceramics (Section 9.6.4). Those doping ions can be introduced as fluorides, oxyfluorides, oxides or even chlorides. According to a study on GeO2:PbO:PbF2þErF3 glass-ceramics by M. Mortier et al. [20], the anionic environment of Er3þ plays an important role on the glass-ceramics formation. For the purpose of this study, Er3þ was introduced as ErF3, ErOF, Er2O3 and ErCl3. It was first noticed that not all glasses could lead to the partial crystallization of PbF2, whichever thermal treatment was applied, perhaps due to different solubility of the various compounds. Second, it was assumed that the first anionic shell surrounding Er3þ is conserved through the melting process, explaining the variation of lifetime of the fluorescing 4I13/2 level in glasses doped with the different erbium precursors, as reported in Figure 9.7. The following section of this chapter will present the most commonly studied oxyfluoride glass-ceramic systems. The synthesis method is reported for each system. However, as

280

Functionalized Inorganic Fluorides (a)

(b)

<100>

Pb2+ Ln3+ F–

<111>

Figure 9.6 (a) Representation of a MF2 unit cell (M ¼ Ca, Cd, Pb, Sr). The black circles represent M2þ cations. For the sake of clarity, F anions, which sit at the corner of each small cube, are not represented on this figure. (b) Representation of the unit cell when doped with Ln3þ. The interstitial fluorine ion compensating the positive charge can for example be situated along the <100> direction or along the <111>

ErF3 ErCl3 ErOF3

4.5

τ (4I13/2)(ms)

4.0

Er2O3

3.5 3.0 2.5 2.0 1

2

4 3 Er ions (mol%)

5

6

Figure 9.7 Variation of the 4I13/2 level lifetime of Er3þ of 50GeO2:40PbO:10PbF2 glasses according to the nature of the erbium precursors (ErF3, ErCl3, ErOF and Er2O3) and for various Er3þ concentrations. Reproduction from reference 20. Copyright (2003) Elsevier Ltd

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the exact synthesis conditions are very specific to the composition of the material, and in particular, to the quantity of doping ions and of fluoride phase [12], the values reported are average values. The ability of each system to incorporate the doping ions into their crystallites is closely examined as it plays a fundamental role in their optical properties. 9.3.1

Glass-Ceramics with CaF2 as their Crystalline Phase

Among the glass-ceramics studied with CaF2 as the crystalline phase, the most common system has the following composition: 45SiO2:20Al2O3:10CaO:25CaF2. The main synthesis method of this material is by partial glass devitrification. The alumino-silica glass is made at 1400 C. To crystallize the calcium fluoride phase, the as-made glass is annealed between 2 and 4 hours at 450 C, leading to transparent glass-ceramics with crystallites of an average size of 10–20 nm [21, 22]. These glass-ceramics have commonly been doped with Eu2þ, Nd3þ and Er3þ. In the case of Eu2þ, time-resolved spectroscopic studies have shown that Eu2þ ions are at least partially incorporated inside the CaF2 crystallites. Similarly, Er3þ ions are incorporated in the crystallites, forming a solid solution Ca1xErxF2þx. However, the incorporation of Nd3þ in the CaF2 crystallites seems weaker [23]. 9.3.2

Glass-Ceramics with b-PbF2 as their Crystalline Phase

Some germanate and silicate glass-ceramics have been synthesized with -PbF2 as the crystalline phase. Similarly to the CaF2 system presented before, these glass-ceramics are made by glass devitrification. Germanate glass-ceramics with composition GeO2:PbF2þErF3þYbF3 were first synthesized by Auzel et al. in 1975 [24]. However, at that time, they were not transparent because of a very high crystallized fraction with micron size crystallites. Later, transparent glass-ceramics with similar compositions (GeO2:PbO:PbF2þErF3) were developed. The initial glass is made by melting GeO2, PbO, PbF2 and LnF3 (Ln ¼ Er [25], Yb [26], ErþYb [26], Er þ Yb þ Ce [27], etc.) powders around 1000 C for 20 to 40 min. In the case of a doping with 2 to 4 % (mol) ErF3, annealing the as-made glass at 360–380 C for 10 hours induces the crystallization of the -PbF2 phase. It leads to a transparent glass-ceramic with Pb1xErxF2þx crystallites of 10 to 20 nm according to the exact composition, as well as the annealing time and temperature. From various studies, it has been proved that the incorporation of Er3þ ions inside the PbF2 network is total if the annealing conditions are carefully chosen [12]. 50SiO2:50PbF2 glass-ceramics doped with 4 % (mol) ErF3 or 1 % (mol) EuF3 can also be synthesized [28]. Some optical characterizations evidence the segregation of Er3þ (or Eu3þ) inside -PbF2 nanocrystallites. 9.3.3

Glass-Ceramics with CdF2/PbF2 as their Crystalline Phase

Transparent glass-ceramics, with a molar composition 33SiO2:17Al2O3:22CdF2:28PbF2, doped with ErF3 and YbF3 were the first transparent glass-ceramics studied for their optical

282

Functionalized Inorganic Fluorides

properties, by Wang et al. in 1993 [29]. Since then, various groups have pursued their study [30] and several alternative compositions, such as 40SiO2:30PbF2:30CdF2 [31], have been envisaged. All of these systems are obtained by glass devitrification. The synthesis of the initial glass is performed by melting some polycrystalline powders of alumina, of silica and of the required fluorides, at 1000 C for 20 to 30 minutes and then by a fast quenching. The glass obtained is then annealed at 450 C for a time varying between 30 minutes and 36 hours, according to the exact composition of the glass. This thermal treatment leads to the crystallization of a fluoride phase, described as a solid solution of CdF2-PbF2 by some authors [32] or only a pure PbF2 phase by some others [33]. The obtained fluoride crystallites have a size ranging from 9 to 20 nm. Such glass-ceramics have been doped with various rare-earth ions, such as Er3þ, Eu3þ [34], 3þ Pr [35], Ho3þ [36] and even doped with two different ions Er3þ/Yb3þ, as reported by Wang et al.[29]. Spectroscopic studies show that rare-earth ions are incorporated well inside the crystallites after devitrification. 9.3.4

Glass-Ceramics with LaF3 as their Crystalline Phase

Glass-ceramics with LaF3 as a crystalline phase have been discovered in 1998 by Dejneka [5]. The basic composition is: SiO2, Al2O3, Na2O and LaF3 doped with ErF3, but other compounds have been added, such as AlF3. The synthesis is usually performed by glass devitrification, with an annealing temperature around 650 C. However, some examples of soft chemistry-based synthesis can also be found [3,37]. In both cases, the glass-ceramics obtained contain nanosized LaF3 particles. A good incorporation of the trivalent rare-earth ions in the LaF3 lattice can be expected as the doping rare-earth ions have the same valence state as La3þ. In such system, Goutaland et al. have shown, using optical spectroscopy, that 70 % of Pr3þ ions were incorporated inside the LaF3 crystallites but that only 2 % of Er3þ were incorporated into this fluoride network [38]. This difference can be understood when taking into account the ˚ , r(Pr3þ) ¼ ionic radii r of the different 8-coordinated lanthanide cations (r(La3þ) ¼ 1.32 A 3þ 3þ ˚ ˚ 1.28 A, r(Er ) ¼ 1.14 A). This bad incorporation of Er ions to the LaF3 crystalline network has also been evidenced by other researchers. In particular, due to the study of upconversion processes in those materials, G. C. Jones et al. evaluated the incorporation of Er3þ inside the crystallites to 20 % of the total doping amount [39].

9.4

Thermal Characterization

The most efficient and sensitive tool to characterize a glass that can give rise to a glassceramic is differential thermal analysis (DTA). It allows the different characteristic temperatures of the material that permit prediction of the glass behaviour under thermal treatments to be determined. To ensure the reproducibility of the characteristic temperatures and to compare various samples or experiments, the thermal analysis has to be carried out carefully at a same heating rate, often 10 C/min. In some cases, a series of DTA

Oxyfluoride Transparent Glass Ceramics

283

measurements can be made with different heating rates to determine the process of crystallization and its associated activation energy [1] as developed in Section 9.4.1 with Johnson-Mehl-Avrami or Chen and Ozawa methods. The experiment also has to be done with a similar powder size distribution, typically between 40 and 70 microns on an equal sample mass. Indeed, some crystallization processes are surface-induced and they could occur for lower temperatures because of higher specific surfaces induced by smaller grain size of a powdered glass. The use of different-sized powders is the very principle of Thakur’s method [40] for detecting the nucleating character of species and the surface role in the crystallization as shown in Section 9.4.2. The thermal event that arises at the lowest temperature is the glass transition. It is generally followed by a series of exothermic contributions reflecting the successive crystallizations of the various phases that the glass composition can give. Some exothermic peaks can also be associated to structural phase transitions of some of the crystallised phases. The stability of a glass can be estimated by several criteria [41, 42] that are moreor-less based on the extent of the temperature interval lying between the glass transition temperature, Tg, and the onset of the lowest crystallization peak. A wide temperature interval, e.g. 100 C, corresponds to the possibility of heating a glass above Tg to shape it, or to release strains induced by the synthesis, with a quasi null risk of devitrification. A reduced temperature interval induces a strong risk of crystallization of the glass matrix with any reheating of the glass around Tg. The simplest stability criterion is just the temperature difference between Tg and the onset, or the maximum, of the first exothermic peak. It is probably the most efficient because more sophisticated factors reveal a weak variation with regard to the uncertainty of their value [43]. At higher temperatures, one or several endothermic peaks are observed. They correspond to the melting of the different phases that have previously been crystallized during the heating process of the thermal analysis. The highest one can be seen as the melting point of the glass. Observation of a DTA curve is sufficient to determine the capability of a glass to give rise, or not, to a glass-ceramic. For the transparent glass-ceramics we are dealing with in this chapter, the requirements are stronger than for non-transparent glass-ceramics. In most cases, to get transpareny, the crystallized part of the material must remain clearly of minor volume with regard to the glassy matrix. Such a configuration corresponds to the observation of several and clearly separated crystallization peaks with the one lying at the lowest temperature being of minor integrated intensity when compared to the whole exothermic contributions. It has ideally to lay far enough above Tg to insure a well-controlled crystallization with heating treatment. The DTA curve associated with an ideal situation and corresponding to the previously described glass composition is shown in Figure 9.8. On this curve, the crystallization peak of -PbF2 is lying between the glass transition, Tg, and the crystallization of the oxide matrix. Such a favourable case is generally obtained by varying the composition of the starting glass to shift the various thermal events adequately, i.e. the glass transition and the various crystallization peaks. The choice of the chemical composition is thus of main importance. As a matter of fact, DTA curves of series of various glass compositions have to be measured to determine the influence of the components on the physical changes of the glass. The key to obtaining a partial and controlled crystallization is generally to observe two or more crystallization peaks that are sufficiently separated so that the precipitation of one phase only can be obtained. In the case of oxyfluoride systems, the oxide matrix must offer

Functionalized Inorganic Fluorides Heat Flow (arb. units) exothermic

284

200

Tc(PbGeO3) Tc(β–PbF2)

Tc(PbGe4O9)

Tg

300

400 500 Temperature (°C)

600

700

Figure 9.8 DTA curve of the glass 50GeO2:40PbO:10PbF2þ0.5ErF3þ1YbF3þ0.5CeF3. Reproduced from reference 27. Copyright (2005) American Chemical Society

a stability range wide enough to allow crystallizing only the fluoride phase. It has been observed that a homogeneous nucleation of a phase can be forced by the increase of the proportion of this phase in the material. For instance, PbF2 crystallization peak temperature is strongly lowered by increasing its content from 10 to 20 mol %. This effect can be understood as similar to the oversaturation of a solute in a solution. Also, heterogeneous nucleation can be promoted by the addition of nucleating agent, such as rare-earth fluorides, which have been demonstrated to be efficient for PbF2 crystallization [1,27]. In Figure 9.9, the effect of the erbium fluoride content is evidenced by the dramatic shift of the exothermic peak of PbF2 down to Tg with the increase of ErF3 concentration. An excess of ErF3 can even induce a shift virtually through the glass transition zone and prevent the synthesis of a fully amorphous material during the synthesis of the glass. This situation could correspond, for instance, to the case of 4 mol of ErF3 that should be extrapolated on Figure 9.9. In those conditions, a non-transparent glass-ceramic is directly obtained when the liquid is quenched. In fact, a faster quenching rate, if possible, could theoretically allow a fully amorphous material to be obtained. But, in practice, very fast cooling rates induce severe strains in the glass, which often explodes just after quenching the melt. In a similar way to the addition of ErF3, the increase in PbF2 itself induces the shift in the crystallization peak of PbF2 down to Tg as observed in Figure 9.10. The glass is then less stable with regard to devitrification. The nucleation process becomes difficult to manage safely and crystallites of greater sizes are obtained with an inherent reduction in the transparency of the material. However, the glass can become self-nucleating and either no or very few rare-earth ions can be used to obtain a glass-ceramic [1]. Most attention has to be paid to the nucleating character of ions or species. Indeed, an extensive study has been performed on erbium compounds used as seeds for nucleation in the GeO2:PbO:PbF2 glass system. The following compounds have been compared: ErF3, ErOF, Er2O3, ErCl3. Their content has been varied from 0 to 2 mol % of erbium and the DTA curves have been compared. The DTA curves of glasses containing the same

Oxyfluoride Transparent Glass Ceramics

285

doped by x(mol%) ErF3

*

DTA signal (a.u.)

x=3

*

x=2

x=1

x=0

250

300

350

400

450

500

550

600

650

Temperature (°C)

Figure 9.9 DTA curves of the glass 50GeO2:40PbO:10PbF2 þ xErF3, with x ¼ [0,1,2,3]. Reproduced from reference 40. Copyright (2003) Elsevier Ltd

Temperature (°C)

520 Tp1

480 Tp (β –PbF2)

440 8

10

12

14 16 y PbF2

18

20

Figure 9.10 Evolution of the characteristic temperatures measured on the DTA curves of the non doped glasses (50GeO2:[50-y]PbOþyPbF2) versus y. Tp(b-PbF2) corresponds to the crystallization temperature of b-PbF2, and Tp1 is associated with the crystallization of the oxide glassy matrix. Reproduced from reference 1. Copyright (2000) Springer Science þ Business Media

proportion of Er3þ ions but introduced as different compounds mostly differ when the erbium compound level is increased up to 2 mol %. The effect is observed for the various crystallization peaks and mainly for the PbF2 peak that is only observed with ErCl3 and ErF3 doping (Figure 9.11). The crystallization conditions are strongly affected too and

286

Functionalized Inorganic Fluorides

DTA signal (a.u.)

2mol% ErCl3

2mol% ErF3

2mol% ErOF

1mol% Er2O3 undoped

280

320

360

400 440 480 Temperature (°C)

520

560

Figure 9.11 DTA curves of the glass 50GeO2:40PbO:10PbF2, with various erbium compounds used as doping species. The arrow points out the crystallization peak of the PbF2 phase. Reproduced from reference 20. Copyright (2003) Elsevier Ltd

only the glass doped with erbium fluoride underlies an easy nucleation and growth of PbF2 particles as expected by the shape and position of its exothermic contribution. After a long time and high-temperature melting in an oxidizing atmosphere, the initial coordination sphere of the erbium ions is mainly preserved in the glass even with an initial ionic bonding such as fluoride or chloride compounds in a predominantly oxide matrix. This effect should be employed to modify the environment of rare-earth ions in a glass and subsequently the optical properties of the glass (Figure 9.7). The significant difference in the solubility of the various rare-earth compounds should also be used to improve the introduction of the rare-earth ions in glasses. In oxyfluoride systems, too, it is possible to modify the composition of the fluoride phase in order to lower its crystallization temperature. For instance, the phase diagram PbF2-CdF2 has been explored by several authors [29, 30, 31]. The existence of a solid solution permits to lower the precipitation temperature without external nucleating agent or excessive content of the phase itself inducing homogeneous nucleation. The segregation of the rare-earth ions in the crystallized phase is of the greatest importance to give improved optical properties to glass-ceramics when compared to glass. The partition of the rare-earth ions, or other active species such as 3d metal ions, in the crystal phase, confers optical properties similar to those of the corresponding single-crystal. However, the success in the partial devitrification of a glass does not induce the automatic segregation of the RE or 3d ions in the crystal phase. The problem of segregation is bound to the role of the ions in the nucleation process. When acting as seeds for nucleation, the ions

Oxyfluoride Transparent Glass Ceramics

287

have the biggest chance to segregate in the nucleated phase. But, as it has been demonstrated with several different erbium compounds, the ion cannot be considered without its anionic environment. In this frame, one can understand that in several glass-ceramic systems for which nucleation is enhanced by the solid solution effect, the rare-earth ions are only weakly partitioned into the crystal phase. Most of the papers dealing with PbF2-CdF2 report strong but partial segregation [28, 29] and sometimes weaker segregation. The ion to partition has to play a role in the crystallization process to ensure its perfect segregation. Indeed, a system in which the nucleation tendency has a homogeneous nature does not need external elements to crystallize. There, the doping ion is weakly incorporated in the crystal phase, unlike in the case of compounds that are stable with regard to the crystallization processes and that need external seeds. The various glass-ceramic systems reported in the literature tend to support such an assumption. A relatively weak PbF2 content in (GeO2:PbO:PbF2) GPF system, i.e. 10mol % PbF2, ensures a high segregation rate of the rare-earth ions in the crystallites (Yb, Er, Ce) because of a broad Tx-Tg interval. On the contrary, compounds with high PbF2 content were not showing segregation of the rare-earth ions into the crystal phase [24] but only a statistical repartition. Then, they are equally present in the whole material and the physical properties are a superposition of the physical properties of the ion in both phases. Also, LaF3 containing glass-ceramic systems have been shown to incorporate a limited part of the rare-earth ions into the crystal phase [38]. The key role of impurity ions has been extensively observed through DTA measurements and crystallization process observations. It has also been studied through molecular dynamics (MD) simulations of the devitrification process of a lead fluoride glass doped with Er3þ ions. The first steps of the formation of Pb1xErxF2þx crystallites were examined using Buckingham-type potential between ions [44, 45]. The total enthalpy, the radial distribution functions and the diffracted intensities of systems containing different amounts of Er3þ ions were modelled. The nucleating role of Er3þ ions has been clearly evidenced through the lowering of the devitrification temperature of PbF2, as shown in Figure 9.12, in good agreement with the experimental results. –17 800

Enthalpy (eV)

–17 850

PF

–17 900 –17 950

PFE1

–18 000 –18 050

PFE4

–18 100 300 320 340 360 380 400 420 440 460 480 500 Temperature (K)

Figure 9.12 Evolution of the total energy during the devitrification process for various cells containing no Er ions (PF), 1 Er (PF1), 4 Er (PF4). Reproduced from reference 44. Copyright (2007) American Institute of Physics

288

9.4.1

Functionalized Inorganic Fluorides

Kinetics of Phase-change/Devitrification

The kinetics of phase change can be studied through the analysis of the thermal measurements by a Johnson-Mehl-Avrami model. Such a study has been carried out on a (50GeO2:40PbO:10PbF2þ3ErF3) glass [25]. A nonisothermal method, with different values of heating rates ( ¼ 5, 6, 7 and 10 C/min), has been used to determine the Avrami parameter. The main results of DTA measurements with different heating laws are: Tp, the temperature corresponding to the maximum of the precipitation peak; the crystallized fraction w(T ) at a given temperature T determined by dividing the integrated intensity under the DTA curve from the beginning of the crystallization peak up to the given temperature T and up to the end of the crystallisation peak. The Avrami parameter, n, or reaction order, is the slope of a linear law fitting ln(ln(1w)) versus ln(). The following values have been obtained for the different fixed temperatures (T ¼ 390, 392, 394 and 396 C) respectively (n ¼ 2.55, 2.24, 2.08, 1.89). Such values of the Avrami parameter correspond to a three-dimensional crystal growth controlled by the diffusion and with a decreasing nucleation rate [46]. They describe in a satisfactory way the devitrification process in which the crystallites begin to grow and then slow down their growth rate with time as shown by the x-ray estimated sizes of the crystallites versus time ([1] and [47]). This behaviour is due to the progressive depletion of lead fluoride around the growing crystallites, involving longer diffusion distances. This is also characteristic of a multi-component glass including many different noncrystalline phases forming a complex heterogeneous system. The decreasing nucleation rate reflects the progressive disappearance of the smallest crystallites in benefit of the largest ones during the Ostwald ripening [48] occurring during the growth period. The crystallization apparent activation energy of the glass was evaluated using the Chen [49] and Ozawa [50] methods. A linear fit of ln ðTp2 =Þ versus (1000/Tp) for the Ozawa method gives a value of 228 kJ/mol. The Chen method, using always a linear fit, with ln() versus (1000/Tp), gives an energy of 239 kJ/mol. For a stable glass like ZBLAN, the activation energy is of 196 kJ/mol [51]. A family of unstable fluoride glasses [52] inclined to spinodal decomposition, has an activation energy of about 345 kJ/mol. So, for a GeO2:PbO:PbF2 glass, a value of about 228 kJ/mol, lying between the values of stable and unstable glasses, reflects the tendency to partial devitrification.

9.4.2

Thakur’s Method

To confirm the nucleating agent role of ErF3, Thakur experiments [53] have been carried out on the GeO2:PbO:PbF2 glass system. This method consists of making a series of DTA experiments on a glass, varying the content of the supposed nucleating agent. For each composition, the DTA experiment is done on a coarse powder and on fine powder, the size being kept constant with the same amount of glass and a similar thermal history through 0 the whole series (Figure 9.13). For each crystallized phase, one can calculate DTm , corresponding to the difference between the crystallization temperature of coarse and fine particles of the glass. It quantitatively represents the change in the activation energy 0 (DE) due to a compositional change. DTm can then be plotted versus the concentration of

DTA signal (a.u.)

Oxyfluoride Transparent Glass Ceramics

very fine

*

fine

*

coarse

320

289

*

400

480

560

640

720

Temperature (°C)

Figure 9.13 DTA curve of the 50GeO2:40PbO:10PbF2þ2ErF3 glass versus the grain size of the powder; coarse: size > 100 mm, fine: size < 71 mm, very fine: 20 < size < 40 mm. The lines are just guides for the eye. Reproduced from reference 40. Copyright (2003) Elsevier Ltd

the doping species. This difference should tend to zero for concentrations where the nucleation will be the most efficient [53]. 0 The values of DTm are obtained by making the difference between the position of the peaks of the two lower curves plotted on Figure 9.13. The results for the various ErF3 concentrations, 0–4 mol %, are presented in Figure 9.14. For the peak associated to PbF2 0 crystallization, DTm shifts down to zero when the erbium content is increased, whereas it remains >10 C for the two other peaks corresponding to the crystallization of PbGeO3 and PbGe4O9. Following the Thakur criterion [53], ErF3 acts as an efficient volume nucleating agent for PbF2 and not for PbGeO3 and PbGe4O9. In turn, it has been shown 0 on the Figure 9.14 and with the values of DTm 10 C that PbGeO3 and PbGe4O9 are very sensitive to surface-enhanced crystallization.

9.5

Morphology of the Separated Phases

Each glass-ceramic sample is first characterized by X-ray diffraction to determine the nature of the crystallized phase and to estimate the average crystal size and the crystallized fraction. Also, the lattice parameter is a powerful tool, through the Vegard law, to measure the rare-earth content of the crystallites when a solid solution exists. In every case, XRD gives only average information on the sample. The direct observation of the morphology of the composite system is done with transmission electron microscopy. The various observation modes allow a complete characterization. A lower resolution micrograph provides information about the size distribution. In high resolution modes, the atomic planes are directly observed and the interatomic distance is measurable. The selected area electron

290

Functionalized Inorganic Fluorides 40 PbGe4O9 PbGeO3

ΔT’m(°C)

30

PbF2 20

10

0 0

1

2

3

4

ErF3 (mol%) 0

Figure 9.14 DTm (C) difference between the position of the exothermic peaks of the coarse and fine powder on the DTA curves versus the ErF3 addition in the 50GeO2:40PbO:10PbF2 glass. Reproduced from reference 40. Copyright (2003) Elsevier Ltd

diffraction, SAED, enables the determination of the crystal phase of each particle. Dark field images also allow a clear separation of somewhat agglomerated particles. X microanalysis permits the determination of the local composition of a particle and of the surrounding glass matrix. The ratio of rare-earth ions segregated in the particles can then be accurately determined. Some examples of informative TEM images are shown in Figures 9.15 and 9.16. The nucleating character of the different rare-earth fluorides is strongly different and dramatically affects the morphology of the separated phases. For instance, the respective

300 Å

Figure 9.15 Transmission electron micrograph of a 50GeO2:40PbO:10PbF2þ3ErF3 glass showing crystallites after treatment 10 h at 360 C. Reproduced from reference 1. Copyright (2000) Springer Science þ Business Media

Oxyfluoride Transparent Glass Ceramics

291

b

d

a

40 Å

c

Figure 9.16 HRTEM image of a 50GeO2:40PbO:10PbF2þ3ErF3 glass-ceramic. Many crystallites showing different atomic planes are observed: (a) (111), (b) (200), (c) (220), (d) crystallite in zone axis. Reproduced from reference 1. Copyright (2000) Springer Science þ Business Media

efficiencies of Yb, Er and Ce fluorides have been compared. The efficiency can be logically associated with the tendency to lower the crystallization temperature and to lower the crystal size but to increase the crystal number. Figure 9.17 allows such an easy comparison. The size and number of crystallites is not the only morphological aspect because the shape of the crystallites is also strongly modified. Figures 9.15 and 9.16 show the aspect of nanoparticles of ErF3 doped GeO2:PbO:PbF2 glasses. Figure 9.18 demonstrates the irregular shape of the particles of PbF2 nucleated with YbF3 doping. Less numerous but bigger than with ErF3, the particles with YbF3 are also grouped together with the same crystallographic orientation on intermediate range. The determination of the crystallized fraction of the glass-ceramic system is a quite delicate problem [1]. The XRD data offer an overall view of the two phases, as shown, for instance, in Figure 9.5. The broad lines of the glass matrix and the sharp lines of the crystallites are superimposed. However, the ratio of the integrated intensities diffracted by the two separated phases is not directly usable when attempting to reach conclusions about the ratio of the phase volumes because the XRD intensity arises both from the atomic diffusion factors and the structure factor of the phase. So, the intensity of a line is not only proportional to the diffracting volume of the phase but also integrates the atomic diffusion factors of the different atoms that constitute the phase and also their periodic, or otherwise, arrangement. Such a ratio between the broad line of the glass and the sharp lines of the crystallites can only be used inside a defined system in order to compare the crystallization

292

Functionalized Inorganic Fluorides 420 (Ln = Er)

Tc(β–PbF2) (°C)

410

(Ln = Yb)

400

390 (Ln = Ce) 380 8

16 12 Crystallites size (nm)

20

Figure 9.17 Evolution of b-PbF2 crystallization temperature and of the crystallite size according to the doping nature, for the following composition: 50GeO2:40PbO:10PbF2þ0.5ErF3þ 1YbF3þ0.5LnF3 (Ln ¼ Ce,Er,Yb). Reproduced from reference 27. Copyright (2005) American Chemical Society

Figure 9.18 TEM image of 50GeO2:42PbO:8PbF2þ3YbF3 glass-ceramic. Reproduced from reference 26. Copyright (2005) Materials Research Society

degree as a function of a thermal treatment or a weak modification of the composition [26]. To obtain an absolute value of the ratio of the phase volumes necessitates preparing standard samples of pure glass and the pure crystal phase before measuring glass-ceramic samples. Also, the DTA crystallization peak surface can be used to estimate the quantity of the crystallized phase, mainly for the weakest ratio at the early stages [12]. The Raman diffusion curves can also give such information [1].

Oxyfluoride Transparent Glass Ceramics

9.6

293

Optical Properties of Glass-Ceramics

We know that, in the best cases, Ln3þ ions are segregated in the crystallites. In this section we will discuss and compare the optical properties (absorption, emission, lifetime) of glass-ceramics, glasses and single-crystals. Throughout the section, we will see the consequences of the segregation of Ln3þ in fluoride crystallites on the glass-ceramic optical properties. 9.6.1

Influence of the Devitrification on the Spectroscopic Properties of Ln3þ

As was mentioned earlier, the devitrification process results in a change of the environment of the lanthanide ions (or, at least, some of them), from being in an oxyfluoride glass to being segregated in fluoride crystallites. These structural and chemical changes bring about changes in the spectroscopic properties of Ln3þ. In the following section, we will compare the optical properties of glasses and their corresponding glass-ceramics. 9.6.1.1

Reduction of the Inhomogeneous Optical Linewidth and Increase of the Absorption/Emission Cross-Section

In Ln3þ-doped glasses, Ln3þ ions are randomly distributed and occupy many different sites. The doping ions are subject to various environments and thus, different crystal-field perturbations. As a result, they have slightly different energy-level splittings and their optical transitions occur at slightly different wavelengths, leading to rather broad absorption/emission bands. On the contrary, in single-crystals, Ln3þ ions have well defined sites and exhibit absorption/emission with narrow inhomogeneous linewidth. In the glass-ceramic systems where the Ln3þ ions are well incorporated into the crystalline nanoparticles, the linewidth of an optical transition is narrower than it is in glasses. It results from the segregation of Ln3þ in the nanoparticles which provide them a crystalline environment. This is illustrated here with the 50GeO2:40PbO:10PbF2 system doped with ErF3, where it was proved that all Er3þ were incorporated into the PbF2 nanocrystals [25]. As shown on Figure 9.19, the absorption band centred at 540 nm, corresponding to the 4I15/2 ! 4S3/2 transition of Er3þ is much narrower in the glass-ceramic than in the as-melted glass. In this particular example (Figure 9.19), the reduction of the inhomogeneous linewidth induces an increase of the maximum cross-sections in the glass-ceramic. Similarly for the emission, in the SiO2:Al2O3:CdF2:PbF2 system doped with NdF3, an increase of the stimulated emission cross-section by a factor of 1.5 in the glass-ceramic compared to the glass is reported for the 4F3/2 ! 4I11/2 transition [54]. Moreover, Quimby et al. experimentally demonstrated that this change of environment from amorphous to crystalline increases the quantum efficiency of Pr3þ at 1.3 mm [55]. It suggests that Ln3þ-doped glassceramics have real potential for optical amplification. However, an increase of the cross-section or quantum efficiency after the devitrification process is not always observed as the change from an oxide environment to a fluoride environment tends to decrease the oscillator strength [56]. Also, crystal-field strength is weaker than Stark level splitting in a fluoride environment compared to an oxide one.

294

Functionalized Inorganic Fluorides

T = 300 K

0.32

heat treated

σa (10–20 cm2)

0.24

0.16

0.08

as melted

0.00 535

540

550 545 Wavelength (nm)

555

560

Figure 9.19 (a) Absorption cross-section, in 1020cm2, of the 4I15/2 ! 4S3/2 transition of Er3þ in an as-melted GeO2:PbO:PbF2 glass and in the corresponding glass ceramic (resulting from a 360 C treatment for 10 hours), at 300 K. Reproduced from reference 25. Copyright (1999) Elsevier Ltd

9.6.1.2

Lengthening of the Excited State Lifetime

Not only does the devitrification have an effect on the Ln3þ optical linewidth and the Ln3þ transition intensity but also it changes the Ln3þ excited state lifetime. Indeed, in glassceramics, the excited state lifetime is usually much longer than in the as-melted glasses. In glass-ceramics, Ln3þ ions are in a fluoride environment, with phonons of much lower energy than in an oxide or oxyfluoride environment. Table 9.1 summarizes the phonon energy cut-off for the different fluoride compounds. For comparison, the phonon energy cut-off in oxides is between 1100 and 1400 cm1 [57]. In a low phonon energy environment, fast nonradiative de-excitations are less likely, inducing an increase of the optical lifetime. To illustrate this lengthening of optical lifetimes in glass-ceramics, the GeO2:PbO:PbF2 system doped with Pr3þ is considered [58]. The fluorescence decay curves from the1D2 level of Pr3þ are reported in Figure 9.20. In the as-melted glass, the lifetime of the 1D2 level was measured to be 4.1 ms, whereas it is of 109 ms in the glass-ceramic resulting from a 395 C treatment for 5 hours. There is an increase by a factor of 26 between the 1D2 excited state lifetimes of Pr3þ from the glass to the glass-ceramic due to the change from an oxide to a fluoride environment. Glass devitrification induces changes in the optical properties of the materials. First, it provides a crystalline environment for the Ln3þ ions, inducing a reduction of their optical

Oxyfluoride Transparent Glass Ceramics

295

Table 9.1 Phonon energy cut-off of different fluoride phases used in the synthesis of oxyfluoride glass-ceramics.

h! max (cm1)

CdF2

-PbF2

CaF2

LaF3

384

337

474

350

1D 2

395/5h

τ = 109 μ s as melted τ = 4.1 μ s

0

50

100 150 time (μ s)

200

250

Figure 9.20 Effect of the devitrification on the emission decay curves from the 1D2 level of Pr3þ in the GeO2:PbO:PbF2þPrF3 system. Reproduced from reference 58. Copyright (2004) Elsevier Ltd

linewidth. It also provides them with a fluoride environment, where the nonradiative deexcitations are less favourable, increasing the excited state lifetime. In addition, the segregation of Ln3þ into the fluoride crystallites induces a strong local Ln3þ concentration inside the crystallites, which also affects the optical properties of glass-ceramics as it is discussed hereafter. 9.6.2

Effect of High Local Ln3þ Concentration in Crystallites

Efficient glass devitrification induces the segregation of Ln3þ into the fluoride phase, which usually represents between 5 to 20 % (mol) of the material. The local Ln3þ concentration in the crystallites of the glass-ceramics is between 5 and 20 times higher than in the initial glass. This high local concentration in the crystallites favours Ln3þLn3þ interactions and has an effect on the optical properties, as it is described below. 9.6.2.1

Influence on the Excited State Lifetime

Even if the global concentration remains reasonably low (1 %), the excited state lifetime of Ln3þ in glass-ceramics decreases as the Ln3þ concentration increases. Lifetime measurements were, for instance, performed on transparent oxyfluoride glassceramics with Pr3þ:LaF3 crystallites, doped with different concentrations [59]. When the Ln3þ concentration increases from 0.01 % to 1 %, the decay time is drastically shortened (Figure 9.21), as a result of Ln3þLn3þ interactions. There is a tradeoff in the lifetime

296

Functionalized Inorganic Fluorides

value between the change of environment (which lengthens the lifetime) and the effect of the segregation, which makes the lifetime shorter. Moreover, the glass-ceramic fluorescence decay curves do not increase exponentially over a short time (above 0.5 % in the example reported in Figure 9.21). This can be explained by fast energy transfer between Ln3þLn3þ ions [60]. 1 Emission Intensity (arb.units)

Emission at 600.5 nm

0.1 1.0%

0.01%

0.5%

0.01 0

50

100

150

200

250

Time (μ s)

Figure 9.21 Fluorescence decay curves of glass-ceramics doped with 0.01 %, 0.5 % and 1 % of Pr3þ. Reproduced from reference 59. Copyright (2003) Elsevier Ltd

Strong Ln3þLn3þ interactions favour energy transfer processes, which can lead to concentration quenching, even in materials where the global doping concentration of the material is rather low. They also favour upconversion processes, as illustrated in the following section. 9.6.2.2

Upconversion Fluorescence

Upconversion processes, which convert low energy photons to higher energy photons, result from either photon addition by energy transfer (APTE) or by absorption in the excited state (ESA) [61]. To illustrate the difference in intensity of the upconversion emission in as-melted glasses and glass-ceramics, two examples are given: the 50GeO2:40PbO:10PbF2 system doped with 2 %ErF3 and the 53SiO2:27Al2O3:11Na2O: 1Al2F6:7La2F6 system doped with 0.07 % Er2F6 [62]. Figure 9.22(a) shows the upconversion spectra of Er3þ in an as-melted germanate glass and in a glass-ceramic resulting from a heat treatment at 395 C for 10 hours. Both materials were excited in the infra-red, at 980 nm, and their emission was monitored in the visible range. The upconversion emission is 10 times more intense in the glass ceramic than in the glass. In glass-ceramics, upconversion processes are favoured for two reasons: • The high local Er3þ concentration. Considering a total segregation of Ln3þ into the PbF2 crystallites, the local Er3þ concentration is 20 %. The Er3þEr3þ distance is thus short, favouring the upconversion processes.

Oxyfluoride Transparent Glass Ceramics

297

• The low phonon energy environment, which reduces nonradiative transfers and lengthens excited state lifetimes. In Figure 9.22(b), the evolution of the upconversion intensity is reported as a function of the glass ceramic annealing temperature. Whereas no upconversion is observed in the glass, up-conversion signals increase in intensity as glass-ceramics are annealed at higher temperatures (or during a longer duration [63]). The better the Ln3þ incorporation inside the crystallites, the stronger the upconversion fluorescence. (b) 4

Intensity (a.u)

Intensity (arb.unit)

(a)

10 hr-395°C

2

S3/2

H11/2

4

I15/2

4

λex = 970 nm

I15/2 4

750°C-12hr

F9/2

4

I15/2

700°C-12hr 650°C-12hr

as-made

450

500

600°C-12hr

550 600 Wavelength (nm)

650

700

550°C-12hr as made

400

500 600 Wavelength (nm)

700

Figure 9.22 (a) Upconversion fluorescence spectra of the 50GeO2:40PbO:10PbF2 glass doped with 2%ErF3 and the corresponding glass-ceramics resulting from a 10-hour heat-treatment at 395C. (b) Upconversion fluorescence spectra of the 53SiO2:27Al2O3:11Na2O:1Al2F6:7La2F6:0.07Er2F6 system after various thermal treatments. Reproduced from reference 62. Copyright (2002) Elsevier Ltd

9.6.3

Comparison of the Optical Properties of Glass-Ceramics and Single-Crystals

As previously described, the optical properties of glass-ceramics are rather different from those of the precursor glasses. In this paragraph, a comparison is made between the spectroscopic characteristics of glass-ceramics and those of single-crystals. As very few comparisons of this kind have been reported in the literature, we will only consider the comparison between a -PbF2 single-crystal and a GeO2:PbO:PbF2 glassceramic system with ErF3 crystallites, whose crystalline phase is -PbF2. This glassceramic system has been fully characterized and it was shown that all the Er3þ ions were incorporated inside the PbF2 crystallites [12]. 9.6.3.1

Absorption Spectra

The GeO2:PbO:PbF2 glass-ceramic absorption spectrum is compared to that of -PbF2 single-crystal in Figure 9.23. The glass ceramic, containing 10 %PbF2, is doped with

298

Functionalized Inorganic Fluorides

2 %ErF3, leading to a crystallite composition of Pb0.82Er0.18F2.18 [12]. The single-crystal, with the Pb0.8Er0.2F2.2 composition, was grown by the Bridgman method [12]. The absorption line shape of Er3þ in the glass-ceramic and in the single-crystal is very similar, suggesting that, in both materials, Er3þ ions occupy similar sites. The absorption line of Er3þ in the glass-ceramic is slightly broader than that of Er3þ in the single-crystal, which could be explained by the fact that, in glass-ceramics, Er3þ ions are in nanosized particles and can be perturbed by the neighbouring glassy matrix. The absorption cross-section was found to be higher in the glass ceramic than in the single-crystal. In the nanosized crystallites, some Ln3þ ions are close to the interface with the oxide matrix and are likely to have high oscillator strength [64].

2,0 glass-ceramic

σ abs (10–20 cm2)

1,5

1,0

0,5

0,0

single-crystal 536

538

540

542

544

546

λ (nm) Figure 9.23 Absorption cross-section of the 4I15/2 ! 4S3/2 transition of Er3þ in a 50GeO2: 40PbO:10PbF2 glass ceramic and a b-PbF2 single-crystal

9.6.3.2

Fluorescence Decays

A comparison was established between the lifetimes of Er3þ in the glass ceramic and in the Pb0.8Er0.2F2.2 bulk single-crystal [12]. In the glass ceramic, the 4I13/2 fluorescence lifetime of Er3þ is measured as 2.1 – 0.1 ]ms, whereas it is 2.5 – 0.1 ms in the single-crystal. The shorter lifetime in the glass ceramic is probably due to the effect of the oxide matrix incorporating the nanoparticles. Indeed, several studies have proved that the oxide glassy matrix interacted with the rare-earth ions situated inside the nanosized crystallites and influenced their spectroscopic properties [65, 66]. Indeed, those Er3þ ions close to the nanocrystallite/glass interface are in distorted sites. As the distortions lower the symmetry, this could result in an increase in the electric dipole transition probability and consequently decrease the radiative lifetime. Moreover, those Er3þ ions close to the surface of the crystallites can be sensitive to the presence of oxide ions in their coordination polyhedron, inducing multiphonon nonradiative contribution to the Er3þ de-excitation and lowering the lifetime.

Oxyfluoride Transparent Glass Ceramics

9.6.4

299

Multi-doped Glass-Ceramics

Codoping a material with several different doping ions is a well known and efficient method to improve its absorption and emission properties [67]. As mentioned previously, in glass-ceramics, energy transfers between Ln3þ ions are very efficient because of high local concentration in crystallites and long Ln3þ excited state lifetimes. Taking advantage of this high energy-transfer efficiency, different doping ions were introduced in glassceramics to improve their optical properties. 9.6.4.1

Co-doping with Two Different Types of Ln3þ Ions

Various studies report the optical properties of glass-ceramics doped with two different types of Ln3þ ions: Er3þ/Yb3þ in the SiO2:Al2O3:CaF2 glass-ceramic [68], Er3þ/Yb3þ in SiO2:LaF3 glass-ceramic [69], Nd3þ/Yb3þ in SiO2:Al2O3:CdF2:PbF2:YF3 glass-ceramic [70], Tm3þ/Yb3þ in GeO2:WO3:PbF2 glass-ceramic [71]. Although each type of lanthanide may have specific nucleation efficiency (Figure 9.7) and solubility with respect to the fluoride phase of the glass ceramic, it was shown that it is possible to segregate simultaneously at a high rate two types of Ln3þ inside the same crystallite [26, 72]. As a consequence, very efficient energy transfers between the two types of Ln3þ are observed, as in the case of an Er3þ/Yb3þ-codoped glass-ceramic [68], where Yb3þ acts a sensitizer due to its high and broad absorption band around 1 mm and Er3þ, the activator, luminesces in the infrared range and/or in the visible range by upconversion. This is illustrated in Figure 9.24, showing that the upconversion emission in an Er3þ/ Yb3þ-doped glass ceramic is much stronger than that in Er3þ/Yb3þ-doped glass. Wang et al. estimated that green and red upconversion fluorescences are respectively two and ten times higher in SiO2:Al2O3:CdF2:PbF2 glass-ceramics doped with ErF3 and YbF3 than in fluoride glasses [73]. 4

Intensity (a.u.)

1200

F9/2

4S 3/2

800

400

0 400

2H 2H 1/2 9/2

500 600 Wavelength (n m)

GCErYb GErYb GEr 700

Figure 9.24 Up-conversion emission spectra of Er3þ in an Er3þ-doped SiO2:Al2O3:CaF2 glass (GEr), in an Er3þ/Yb3þ-doped glass (GErYb) and in a Er3þ/Yb3þ-doped glass ceramic (GCErYb), under a 980 nm excitation. Reproduced from reference 68. Copyright (2006) Institute of Physics Publishing Ltd

300

Functionalized Inorganic Fluorides

9.6.4.2

Er3þ/Yb3þ/Ce3þ -doped Glass-Ceramics

GeO2:PbO:PbF2 glass-ceramics doped with ErF3, YbF3 and CeF3 are another example of a system where very efficient energy transfers between Ln3þ occur [27]. Er3þ-doped glass-ceramics are often studied for their potential applications as optical amplifiers at 1.5 mm, emission corresponding to the 4I13/2 ! 4I15/2 transition of Er3þ. To increase the absorption cross-sections at 980 nm, these glass-ceramics are additionally doped with YbF3. However, as mentioned in the previous paragraph, upconversion processes are very efficient in Er3þ/Yb3þ codoped glass-ceramics [73] and they compete with the 1.5 mm emission. For any application at 1.5 mm one should try to minimize upconversion. The introduction of a third doping ion, Ce3þ, in glass-ceramics, has proved to be a very efficient way of reducing upconversion processes. The upconversion emission drops by a factor of 300 (Figure 25.a) at the benefit of the 1.5 mm emission, which is enhanced by a factor of three (Figure 25.b) [27, 74]. For comparison, in glasses, the upconversion emission only decreases by a factor of three with the addition of CeF3. (a)

(b) with Ce3+

350

Intensity (Arb.Units)

Intensity (Arb.Units)

Glass-ceramic without CeF3

Glass-ceramic with CeF3 (Intensity × 100)

400

450

500

λ (nm)

550

600

1400

without Ce3+

1450

1500

1550

λ (nm)

1600

1650

1700

Figure 9.25 (a) Up-conversion spectra of a GeO2:PbO:PbF2 glass ceramic doped with ErF3 and YbF3 (dotted line) and a GeO2:PbO:PbF2 glass ceramic doped with ErF3, YbF3 amd CeF3 (solid line). The intensity of the Ce3þ-containing compound has been multiplied by 100 for the sake of clarity. (b) Emission at 1.5 mm of these glass-ceramics. In both cases, the excitation was performed at 975 nm. Reproduced from reference 27. Copyright (2005) American Chemical Society

This strong decrease of the upconversion emission in glass-ceramics is related to very efficient energy transfers between Er3þ and Ce3þ, called cross-relaxations. Those crossrelaxation processes enable the very fast depopulation of the 4I11/2 excited state of Er3þ onto the 4I13/2 state. Indeed, as shown on Figure 9.26 which represents the fluorescence decay curves of the 4I13/2 level of Er3þ under a 980 nm excitation, i.e. exciting in the 4I11/2 level of Er3þ, no rise time is observed in the glass ceramic containing CeF3 where a rise time of 360 – 20 ms is measured in the one without CeF3. It indicates that, in the presence of Ce3þ, the lifetime of the 4I11/2 level of Er3þ is drastically reduced, preventing any upconversion processes from occurring. It evidences Ce3þEr3þ interactions, which create new nonradiative de-excitation paths from the 4I11/2 to the 4I13/2 level of Er3þ [27].

Oxyfluoride Transparent Glass Ceramics Glass-ceramic without Ce3+

1

Glass-ceramic with Ce3+

0,1 1,0

0,01

Intensity (Arb. units)

log (Intensity) (Arb.units)

301

(b)

0,8 0,6 0,4 0,2 0,0 0,000

0,000

0,002 Time (s)

0,005

0,004

0,010

(a) 0,015

Time (s)

Figure 9.26 Fluorescence decay curves of the 4I13/2 level of Er3þ in a GeO2:PbO:PbF2 glass ceramic doped with ErF3 and YbF3 (dotted line) and in a GeO2:PbO:PbF2 glass ceramic doped with ErF3, YbF3 and CeF3 (solid line): (a) vertical scale log(Intensity), (b) vertical scale Intensity. Reproduced from reference 27. Copyright (2005) American Chemical Society

9.7

Conclusion

We have described various oxyfluoride systems that give rise to a system with two separated phases by nucleation and growth of a fluoride phase inside an oxide glass. Spinodal processes can occur but corresponding glass compositions are generally avoided because of the poor control of the crystallization. We have shown the deciding role of the various components of the starting glass: oxide matrix, fluoride phase, rare-earth compound. The thermal analysis by DTA has demonstrated its utility for describing the glass: the characteristic temperatures and the nature of the crystallization processes. The decisive role of the compound used to introduce the rare-earth ions has been demonstrated through the capability to provoke both nucleation and segregation. Various rare-earth ions have also been shown to influence greatly the morphology of the separated phases because of a variable nucleating character. Efficient segregation of rare-earth ions inside the fluoride crystal phase induces a dramatic reduction of the inhomogeneous absorption and emission linewidths compared to the starting glass. At the same time, the nonradiative processes that affect the fluorescence intensity and shorten the fluorescence lifetime in the parent glass are strongly reduced thanks to the low phonon frequency of the fluoride neighbouring. The segregation of the rare-earth ions inside a minor phase of limited volume generally induces clustering effects and enhances greatly the probability of energy transfers between ions. An extensive understanding of the causes of nucleation processes allows transparent systems to be prepared with crystal-like optical properties of high potential for many applications of interest in the domain of photonics. Because of their specific ability to favour energy transfers, upconversion with photon addition or downconversion with photon cutting [75], and because of their easy fabrication, shaping and sizeability, efficient

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Functionalized Inorganic Fluorides

setups should be developed for optical converters in photovoltaic applications [76]. That constitutes one of the most promising emerging domains today.

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[47] K. Hirao, K. Tanaka, M. Makita, N. Soga, Prepapration and optical properties of transparent glass ceramics containing -PbF2-Tm3þ, J. Appl. Phys. 78, 3445–3450 (1995). [48] J. Zarzycki, Les verres et l’e´tat vitreux, Ed. Masson, Paris, (1982). [49] H. S. Chen, Method for evaluating viscosities of metallic glasses from rates of thermal transformations, J. Non-Cryst. Solids 27, 257 (1978). [50] T. Ozawa, Kinetics of non-isothermal crystallization, Polymer. 12, 150 (1971). [51] M. Matecki, I. Noiret, J. Lucas, Devitrification of a heavy metal fluoride glass, J. Non-Cryst. Solids, 127, 136 (1991). [52] P. Santa-Cruz, D. Morin, J. Dexpert-Ghys, F. Glass, F. Auzel, New lanthanide-doped fluoride based vitreous materials for laser applications, J. Non-Cryst. Solids, 190, 238–243 (1995). [53] R. L. Thakur, in: L. L. Hench, S. W. Freiman (Eds), Advances in Nucleation and Crystallisation in Glasses, The American Ceramic Society, Columbus, OH, (1971). [54] M. Abril, J. Mendez-Ramos, I. R. Martin, U. R. Rodriguez-Mendoza, V. Lavin, A. DelgadoTorres, V. D. Rodriguez, P. Nunez, A. D. Lozano-Gorrin, Optical properties of Nd3þ ions in oxyfluoride glasses and glass ceramics comparing different preparation methods, J. Appl. Phys. 95, 5271–5279 (2004). [55] R. S. Quimby, P. A. Tick, N. F. Borrelli, L. K. Cornelius, Quantum efficiency of Pr3þ doped transparent glass ceramics, J. Appl. Phys. 83(3), 1649–1653 (1998). [56] J. McDouglas, D. B. Hollis, M. J. P. Payne, Spectroscopic properties of Er3þ in fluorozirconate, germanate, tellurite and phosphate glasses, Phys. Chem. Glasses 37, 73–75 (1996). [57] J. M. F. Van Dijk, M. F. H. Schuurmans, On the non radiative and radiative decay rates and a modified exponential energy gap law for 4f-4f transitions in rare-earth ions, 78(9), 5317 (1983). [58] W. Ryba-Romanowski, G. Dominiak-Dzik, P. Solarz, B. Klimesz, M. Zelechower, Effect of thermal treatment on luminescence and VUV-to-visible conversion in oxyfluoride glass singly doped with praseodymium and thulium, J. Non Cryst. Solids 345 and 346, 391–395 (2004). [59] X. J. Wang, H. R. Zheng, D. Jia, S. H. Huang, R. S. Meltzer, M. J. Dejneka, W. M. Yen, Spectroscopy of different sites in Pr3þ-doped oxyfluoride glass ceramics Microelectronics Journal 34, 549–551 (2003). [60] M. Mortier, P. Goldner, C. Chateau, M. Genotelle, Erbium-doped glass-ceramics: concentration effects on crystal structure and energy transfer between active ions, J. Alloys Comp. 323, 245–249 (2001). [61] F. Auzel, Upconversion and anti-stokes processes with f and d ions in solids, Chem. Rev. 104, 139–173 (2004). [62] S. Tanabe, H. Hayashi, T. Hanada, N. Onodera, Fluorescence properties of Er3þ ions in glass ceramics containing LaF3 nanocrystals, Opt. Mater. 19, 343–349 (2002). [63] Y. Kawamoto, R. Kanno, J. Qui, Upconversion luminescence of Er3þ in transparent SiO2-PbF2ErF3 glass ceramics, J. Mater. Sc. 33, 63–67 (1998). [64] J. McDouglas, D. B. Hollis, M. J. P. Payne, Spectroscopic properties of Er3þ in fluorozirconate, germanate, tellurite and phosphate glasses, Phys. Chem. Glasses 37, 73–75 (1996). [65] R. S. Meltzer, W. M. Yen, H. Zheng, S. P. Feofilov, M. J. Dejneka, B. M. Tissue, H. B. Yuan, Evidence for long-range interactions between rare-earth impurity ions in nanocrystals embedded in amorphous matrices with the two-level systems of the matrix, Phys. Rev. B 64, 100201 (2001). [66] M. A. Flores-Gonzales, G. Ledoux, S. Roux, K. Lebbou, P. Pierriat, O. Tillement, Preparing nanometer scaled Tb-doped Y2O3 luminescent powders by the polyol method, J. Solid State Chem., 178, 989–997 (2005). [67] F. Auzel, Compteur quantique par transferts d’e´nergie entre deux ions de terres rares dans un tungstate mixte et dans un verre, C. R. Acad. Sc. Paris 262, 1016 (1966). [68] X. Qiao, X. Fan, M. Wang, J. L. Adam, X. Zhang, Spectroscopic properties of Er3þ/Yb3þ codoped 50SiO(2)-20Al(2)O(3)-30CaF(2) glass and glass ceramics, J. Phys.: Cond. Matter 18, 6937–6951 (2006). [69] A. Biswas, G. S. Maciel, C. S. Friend, P. N. Prasad, Upconversion properties of a transparent Er3þ-Yb3þ co-doped LaF3-SiO2 glass-ceramics prepared by sol-gel method, J. Non Cryst. Solids, 316, 393–397 (2003).

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10 Sol-Gel Route to Inorganic Fluoride Nanomaterials with Optical Properties Shinobu Fujihara Department of Applied Chemistry, Faculty of Science and Technology, Keio University, 3-14-1 Hiyoshi, Kohoku-ku, Yokohama 223-8522, Japan

10.1

Introduction

A sol-gel method, which is the most significant wet process to prepare solids from liquids, can offer interesting and useful routes for synthesizing inorganic metal fluoride and oxyfluoride materials for applications in optics and photonics. Formerly, heavymetal fluoride glasses were intensively studied for potential use in optical fibres. Now the method is extended to prepare a variety of metal fluorides such as alkaline earth and rare-earth fluorides and oxyfluorides in the form of nanoparticles, thin films and nanocomposites. They can be utilized as antireflective coatings, luminescent materials, VUV materials, IR materials, etc. This chapter describes fundamentals and possible applications of optically useful inorganic fluoride nanomaterials through the sol-gel method.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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Functionalized Inorganic Fluorides

Principles of a Sol-Gel Method Metal Oxide Materials

A sol-gel method is well-known as a technology for synthesizing metal oxide materials, including glasses, ceramics and nanomaterials, starting from chemical solutions [1–3]. Silica (SiO2) gels and glasses are some of the most successful examples of sol-gel-derived materials. They can be obtained in the form of dense or porous bulk monolith, fibres and thin films. Many other crystalline oxide materials can also be produced by a sol-gel method by controlling solution chemistry and heat-treatment procedures. Recently, organic-inorganic hybrid materials have attracted much attention as a new class of functional materials that can enhance characteristics of each component. In a typical synthesis, raw materials are first dissolved in a certain kind of solvents. Then the resultant solutions are dried at moderate temperatures and finally converted into solids at higher temperatures. During a sol-gel processing using metal alkoxides, chemical reactions such as hydrolysis and condensation proceed to form sols and gels. For example, in the sol-gel process of tetramethylorthosilicate (TMOS), the hydrolysis reaction generally occurs as follows SiðOCH3Þ 4 þ nH2 O ! SiðOCH3Þ4 n ðOH Þn þ nCH3 OH The formation of metal complexes is also an important step when metal salts (chlorides, nitrates, sulfates, acetates, etc.) are used as raw materials. Instead of the formation of sols, chemically stable metal complexes are formed in solutions. The sol-gel processing of zinc oxide (ZnO) is a good example [4]. Zinc acetate dihydrate (Zn(CH3COO)22H2O) was dissolved in 2-methoxyethanol, and monoethanolamine (NH2CH2CH2OH) was added to the solution. Aminoethanols such as monoethanolamine and diethanolamine act as bidentate ligands to Zn2þ, making the solution stable against any precipitates. Both oxygen and nitrogen in NR2CH2CH2OH (R¼H or CH2CH2OH) can make bonds with Zn2þ. When the precursor was heated at around 300 C, the decomposition began by having ‘(N,O)Zn’ and ‘(COO)Zn’ bonds in the complex broken. The decomposition is followed by the nucleation and crystal growth of ZnO. This kind of a sol-gel reaction is of prime importance for designing chemical reactions in fluorine-containing sol-gel routes based on pyrolysis of organofluorine species.

10.2.2

Metal Fluoride Materials

Low refractive indices and low phonon energies of metal fluoride compounds are suitable for their use in optical devices. High-purity products of metal fluorides have been produced using ‘dry’ processing to avoid oxygen or other contamination. Because a solgel method is a typical ‘wet’ process, attention should be paid to influences of solution chemistry on the purity of final products. Adoption of a sol-gel method to optical fluoride materials started with inorganic fluoride glasses. At first, the sol-gel processing appeared rather challenging because fluoride gels were not known and little was known about the chemical reaction between fluorinating

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reagents and organometallics [5]. Strongly motivated by the thought that the sol-gel processing could utilize cheaper starting materials of high quality and offer a substantial energy savings in processing, it was intensively studied in the 1990s for the fabrication of heavy-metal fluoride glasses. Melling and Thomson first studied a ZrF4–BaF2–LaF3–AlF3 (ZBLA) glass system with zirconium propoxide (Zr(OC3H7)4), barium ethoxide (Ba(OC2H5)2), lanthanum propoxide (La(OC3H7)3) and aluminum isobutoxide (Al(OCH(CH3)C2H5)3) as starting materials [6]. The metal alkoxides were dissolved in tetrachlorotetrafluoropropane (Freon 214) and added dropwise to a mixture of Freon 214 and bromine trifluoride (BrF3) in an aluminum container to obtain amorphous solid products. Since then, a multiple-step process incorporating the sol-gel synthesis and reactive atmosphere processing has been devised to prepare ZBLA glasses [7]. A porous, monolithic, atomically homogeneous zirconium barium lanthanum aluminum hydrous oxide gel was prepared first starting from pertinent metal alkoxides and hydroxides. The resultant oxide gel was then treated in a reactive atmosphere of hydrofluoric acid (HF) and transformed into fluoride glasses. Later works have attempted to improve the processes as well as optical properties of the fluoride glasses. Another promising approach to sol-gel-derived fluoride materials is based on the use of trifluoroacetic acid. The author’s group has succeeded in the sol-gel preparation of fluoride and oxyfluoride coating films [8,9], oxyfluoride glass-ceramics [10, 11] and oxide/fluoride nanocomposite thin films [12, 13]. It had been difficult to prepare directly oxyfluoride glassceramics by the sol-gel method because of the absence of appropriate fluorine sources. When fluoride ions (F) are added to the silicon alkoxide systems, they work as the most effective catalyst in accelerating the gelation process. The F ions attack alkoxides by nucleophilic substitution, which leads to the formation of five-coordinated silicon [1]. New sol-gel technologies are therefore expected to be developed in the fabrication of many kinds of optically functional fluoride-based materials. Some examples will be presented in the following sections.

10.3

Fluorinating Reagents and Method of Fluorination

In a conventional glass technology, fluoride glasses are produced by batching and melting of starting metal fluoride materials. Fluorine is therefore included in raw materials. In this case, the purity of raw materials must be 99.99 % or higher [14]. Contaminations related to oxygen species as well as transition metals should be minimized for the use in high-quality optical components. Glass batches are melted in a platinum crucible under nitrogen. A reactive atmosphere of gaseous inorganic fluorides such as NF3 and SF6 is also employed to remove traces of oxygen. In the sol-gel processing of fluoride glasses, the fluorination treatment is a more important process to obtain chemically and optically pure materials. Therefore, fluorinating reagents such as BrF3, HF, NF3 and SF3 are used for direct fluorination of oxide precursors. It should be noted that BrF3 is a toxic material and an extremely reactive fluorinating reagent with the potential to cause fire. Konishi et al. [15] prepared amorphous 60ZrO2–30BaO–10LnO1.5 gels (Ln ¼ La, Ce, Pr, Nd or Eu) and fluorinated them by NF3. In the fluorination process, the reaction rates were found to be substantially different with the Ln species, becoming slower in the order of La, Ce,

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Pr, Nd and Eu. Saad and Poulain [5] carried out fluorination of a wet oxide gel containing Zr, Ba, La, Al and Na using anhydrous HF as a fluorinating reagent. Figure 10.1 shows a typical experimental setup for fluorination. The reaction time was approximately 1 h for a 10 g batch at 200 C. In this process, fluorination is implemented on an oxide gel, which still contains a significant amount of water and alcohol. At 200 C, the chemical equilibrium oxide þ HF ! fluoride þ H2 O can be largely shifted to the right side, especially for basic oxides. Metal fluoride nanomaterials can also be prepared by fluorination of oxide gels. The reaction of magnesium and aluminum alkoxides with HF in nonaqueous solvents led to the formation of amorphous or crystalline nanosized MgF2 and AlF3, respectively, with high specific surface areas [16, 17].

Furnance

HF

N2

Soda Lime

Figure 10.1 Diagram of the fluorination set-up made from fluorinated polymers (Teflon) and sintered alumina. Reproduced from reference 5 by permission of Elsevier

Trifluoroacetic acid (TFA; CF3COOH) has been widely used as a nonaqueous solvent for both inorganic and organic compounds [18]. TFA is characterized as a strong acid with a larger electrolytic dissociation constant of 5.9 101 due to an electron-attracting trifluoromethyl (CF3) group. Its boiling point is 72.4 C, which is lower that that of acetic acid (118 C). In terms of the fluorinating reagents, TFA is not a reactive fluorine species at moderate temperatures because of the relatively strong C-F bond. Interestingly, however, certain kinds of metal trifluoroacetates decompose into metal fluorides at elevated temperatures [19, 20]. A general reaction scheme for lanthanide trifluoroacetates proceeds as follows: LnðCF3 COOÞ3 ! LnF3 þ ðCF3 COÞ2 O þ CO2 þ CO Using this chemical phenomenon, heavy-metal fluoride glasses were prepared using trifluoroacetates of zirconium, barium, lanthanum, aluminum and sodium [21]. Figure 10.2 shows thermal gravimetric analysis of the ZBLAN powder in this process. A drastic decrease in weight (47.4%) was observed in the range of 220–300 C, which was attributed to the decomposition from the trifluoroacetate to the solid ZBLAN fluoride.

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Weight %

100 80 60 40 20

0

200

400

600

800

1000

t/ °C

Figure 10.2 Thermal gravimetric analysis of the ZBLAN precursor. Reproduced from reference 21 by permission of Elsevier

Trifluoroacetic acid is miscible with other solvents such as ether, acetone, methanol, isopropanol, n-buthanol, benzene, xylene, water, so it can be included in the sol-gel processing to prepare homogeneous precursor solutions. In general, sol-gel precursor solutions are prepared by dissolving metal alkoxides or acetates in organic solvents mixed with TFA. When the solubility of metal compounds is low, the addition of water is helpful to dissolve them completely. In the solutions, the coordination between TFA ions (CF3COO) and metal ions is made because CF3COO acts as a very weak base. Existing anions such as CH3COO and C2H5O are stronger bases and hence associate with hydrogen ions. Other metal salts such as nitrates, chlorides and sulfates should not be used for the processing because the anions (NO3, Cl and SO42) could disturb the coordination between TFA and metal ions. The precursor solutions or dried gels containing metal trifluoroacetate complexes are subject to heat treatments at temperatures higher than 300 C, which is a typical decomposition temperature to form metal fluorides. Exact decomposition temperatures depend on the kind of metal elements involved [19, 22]. At much higher temperatures, metal oxyfluoride or metal oxide materials are also obtained. However, not all the metal elements form fluoride compounds in this process. At least, metal elements with electronegativity less than 1.5 can form metal fluorides through the decomposition in the air. Alkali, alkaline earth and rare-earth fluorides are therefore successfully obtained. Heattreatment atmospheres sometimes influence the decomposition behaviour of trifluoroacetate precursors [20]. Alkaline earth and rare-earth fluoroalkoxides ([M(OR)n]n’ where M ¼ metal elements and OR ¼ fluoroalkoxo groups; R ¼ CH2CF3, CH(CF3)2, C(CF3)2CH2(OCH2CH2)2OMe, C(CF3)2CH2OCH2CH2N(CH2CH2OMe)2, etc.) were examined as precursors for metal fluorides both in the sol-gel process and in the chemical vapour deposition [23, 24]. The fluoroalkoxides were prepared by the reaction of polyether-substituted flouroalcohols with metal source reagents of Sr(OC3H7)2 or BaH2. It was possible to obtain pure metal fluoride powders by hydrolysis or pyrolysis of the precursors. According to the thermal analysis,

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the rapid loss of weight started at 200 C, and the residual weight reached a steady value at approximately 300 C to form metal fluorides. Fluorosilicic acid (H2SiF6) can possibly be a fluorine source in the sol-gel processing. Glass-ceramics in a SiO2–Al2O3–CaO–CaF2 system were prepared from tetraethylorthosilicate (TEOS), Al(NO3)39H2O, Ca(NO3)24H2O, H2SiF6, ethanol and water [25]. An attempt was made to obtain pure and homogeneous glass-ceramics with controlled microstructure for application to dental and orthopaedic cements. Tetraethylorthosilicate was first hydrolysed in ethanol at room temperature. The previously dissolved aluminum and calcium salts were then added dropwise to the TEOS solution. H2SiF6 was added further to the reaction solution. The fully mixed solution was then heated under reflux with continuous stirring until gelation occurred. The preparation of silicon oxyfluoride materials with high fluorine contents was attempted by direct hydrolysis of fluoroalkoxysilane precursors [26]. Silicon oxyfluoride gels of nominal composition SiO(20.5x)Fx were prepared by hydrolyzing triethoxyfluorosilane (Si(C2H5O)3F). Transparent monolithic gels (Figure 10.3) were

Figure 10.3 Silicon oxyfluoride gel prepared under Ar atmosphere. Reproduced from reference 26 by permission of Springer

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obtained under careful control of hydrolysis and condensation reactions. Maintenance of the Si-F bond during gelling, heating and aging was likely to be achieved by treating the gels under an argon atmosphere. One of the most useful materials containing fluorine is fluorine-doped tin dioxide (FTO). This is utilized as transparent conductive films in a variety of optoelectronic devices. Various types of fluorinated tin complexes were developed to fabricate FTO films and nanopowders by the sol-gel method [27–31]. The precursors including Sn-F bonds (for example, see Figure 10.4) are highly useful in obtaining conductive FTO because the amount of fluorine could remain well above the effective doping level. F

R2

OR1

O Sn

R2

R1 = tAm, iPr, Et R2 = Me, tBu

O

O O

R2

R2 Figure 10.4 Structure in solution of the fluorinated SnF(OR)(acac)2 compounds, where R stands for Et ¼ C2H5, iPr ¼ CH(CH3)2 and tAm ¼ C(CH3)2C2H5. Reproduced from reference 31 by permission of Elsevier

10.4

Control of Shapes and Microstructures

Optical properties of materials greatly depend on their microscopic as well as macroscopic shapes, based on the interaction with light as shown in Figure 10.5. Transparency of materials is determined by absorption and scattering of incident light. Reflection at interface is also an important factor influencing optical performance. To make good use of the excellent optical properties of metal fluorides, their morphology and microstructure should be carefully controlled. In the bulk form, heavy-metal fluoride glasses and CaF2 single crystals are the most important fluoride materials as main optical components. Some other single crystals of complex fluorides such as LiYF4 are also important as laser materials [32]. Recently, LuLiF4:Tm,Ho has received considerable attention, although LuLiF4 has some difficulty with crystal growth. Jing et al. succeeded in growing high-quality LuLiF4:Tm,Ho crystals by the Czochralski method [33]. As such, bulk fluoride materials are fabricated generally by high-temperature processes. In contrast, low-temperature processes have been successfully employed for creating fluorine-containing nanomaterials. BaMgF4 and SrAlF5 powders and thin films were first synthesized by the sol-gel method using TFA as the fluorine source [34]. Damien et al. and Lepoutre et al. synthesized LiGdF4:Eu3þ and LiYF4:Er3þ crystallized powders by the solgel method [35, 36]. They showed that the organic groups were removed after annealing treatments at 550 C in F2 atmosphere. Yi et al. prepared high-quality LiYF4, BaYF5 and

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Refraction

Transmission

Incidence

Pores Precipitates Grain boundaries

Reflection

Scattering

Absorption

Luminescence

Electronic transition

Figure 10.5 Various interactions between materials and light

NaLaF4 nanocrystals by codecomposition of precursors in organic solvents [37]. Trifluoroacetic acid and the reaction temperature were found to be crucial for the formation of NaLaF4 and LiYF4 nanoparticles. LiYF4, BaYF5 and NaLaF4 nanoparticles, codoped with Yb (20 %) and Er (2 %), showed bright upconversion fluorescence upon 980 nm NIR excitation. Eu-doped BaMgF4 nanopowders, which were prepared by the trifluoroacetate-based sol-gel method, were successfully used as precursors for the production of a BaMgAl10O17:Eu2þ blue-emitting phosphor at the low heating temperature of 1300 C. This achievement was due to the higher reactivity of the sol-gel-derived materials as well as the lower melting point of BaMgF4 [38]. The sol-gel processing is a versatile method for depositing coatings on a variety of substrates in an economical manner. Thin film is also a common shape of metal fluorides in optics. Fluoride coatings are useful because of their wide range of wavelengths giving high optical transparency. A well-known example is the use of MgF2 in anti-reflective (AR) coatings. LaF3 is an excellent host crystals for luminescent rare-earth ions. Since optical amplification is possible in such materials, an optical waveguide amplifier based on rare-earth thin films was proposed. Buchal et al. studied photoluminescence at 1.5 mm in Er3þ-doped YF3, LuF3 and LaF3 thin films prepared by vacuum evaporation [39]. The sol-gel method using TFA was then investigated to prepare LaF3 and YF3 thin films [9, 40]. Eu2þ-doped BaMgF4 and BaLiF3 thin films, which were fabricated without using reducing atmosphere, could exhibit strong violet/blue luminescence upon irradiation with UV light [41, 42]. Nanocomposite is an ideal structure of metal fluoride nanocrystals for viable optical applications. The idea originates from designing oxyfluoride glass ceramics where nanometer-scale metal fluoride crystals are dispersed in glass matrices of high transparency. Wang and Ohwaki attempted to prepare transparent oxyfluoride glass ceramics based on

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aluminosilicates, in which PbxCd1xF2 microcrystallites codoped with Yb3þ and Er3þ ions were precipitated, by a conventional melting method and observed enhanced optical properties [43]. Since then, various kinds of oxyfluoride glasses have been prepared by using a similar melt-quench processing because they could offer an economical alternative with substantial performance improvements over fluoride glasses [44–48]. However, it has not been possible to fabricate thin films for applications such as planar waveguides. LaF3 is an ideal host for luminescent rare-earth ions because it has extensive solid solutions with all rare-earth ions, a lower phonon energy (350 cm1) than ZBLAN (580 cm1) and adequate thermal and environmental stability [46]. The first attempt to prepare oxyfluoride glass-ceramics by the sol-gel method was therefore made with a simple SiO2–LaF3 composition [49]. Dry silica gels containing La3þ and trifuoroacetate ions (CF3COO) were prepared from TMOS, methanol, dimethylformamide (DMF), lanthanum acetate, trifuoroacetic acid, water and nitric acid. By heating the gels at temperatures above 300 C, LaF3 microcrystals with a size of 10–30 nm were formed in the silica matrix. Because of the good dispersion, LaF3 did not show crystal growth at temperatures up to 800 C. A large specimen of SiO2–LaF3 glass-ceramics was obtained by carefully controlling a composition of the precursor solution (Figure 10.6) [50]. Colorization, crack formation and fragmentation of gels during drying and heating are greatly dependent on the amount of TFA, which acts both as a fluorine source and an acid. It was possible to suppress fragmentation of the heated gels by using lanthanum trifluoroacetate gels as a precursor in preparing the solutions.

Figure 10.6 Appearance of the SiO2 – LaF3 glass ceramics prepared by the sol-gel method. Reproduced from reference 50 by permission of Springer

Transparent oxyfluoride glass-ceramic thin films were prepared using the sol-gel method [10, 13]. SiO2–LaF3 and SiO2–LaOF glass-ceramics were formed by heating at temperatures of 300–500 C and 600–900 C, respectively. Eu3þ activators were successfully incorporated into oxyfluoride crystals, as evidenced by their characteristic luminescent properties. Films of a SiO2–(Gd,Eu)F3 nanocomposite structure, in which (Gd,Eu)F3 nanocrystals were dispersed in a silica matrix, were also deposited on silica glass substrates at temperatures of 300–500 C.

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The above-mentioned pioneering works have been followed by many excellent works on luminescent glass-ceramics. Biswas et al. prepared transparent 0.1ErF3–0.1YbF3–5LaF3– 94.8SiO2 (mol %) glass-ceramics by the sol-gel method using TFA [51]. A bulk sample could be obtained by heating finally at 1000 C. The rare-earth ions (Er3þ and Yb3þ) were partitioned in the low phonon energy LaF3 nanocrystals with average size of 10–20 nm embedded in silica glass, as seen as black spots in a TEM image (Figure 10.7). As a result, efficient visible upconversion emissions were observed under NIR excitation. Luo et al. studied a crystallization behaviour of -PbF2 in transparent glass-ceramics of xErF3– 10PbF2–90SiO2 (mol %, x ¼ 0, 1, 2) compositions by the sol-gel method using TFA [52]. They evaluated the apparent activation energy for the crystallization of -PbF2 phase to be about 162 kJ/mol and 167 kJ/mol for dried xerogels with x ¼ 0 and 1, respectively. It was revealed that few Er3þ ions contributed to or were incorporated into the -PbF2 lattice; rather they seemed to segregate at the surface of the crystallites and hindered the growth of -PbF2. This result is different from the successful incorporation of the rare-earth ions into the LaF3 nanocrystals. Interestingly, the presence of Pb(NO3)2 was observed at around 200 C in their work, which indicated that there existed a competition between NO3 (added to the solution as the acid HNO3 catalyst) and CF3COO to coordinate with the Pb2þ ion.

Figure 10.7 TEM picture of the oxyfluoride glass-ceramic sample showing crystalline precipitates. The inset shows the electron diffraction pattern obtained from the precipitated region. Reproduced from reference 51 by permission of Elsevier

Ribeiro et al. attempted to fabricate Eu3þ and Tm3þ-doped SiO2–LaF3 glass-ceramics as waveguides [53]. Actually, thin films obtained by dip-coating on SiO2 substrates and treated at 300, 500 and 800 C displayed guided modes in the visible and infrared regions. YF3–SiO2 [54], Er3þ-doped BaF2–SiO2 [55], Er3þ-doped SrF2–SiO2 [56], Er3þ-doped

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CaF2–SiO2 [57], Eu3þ- or Er3þ-doped SiO2–LaF3 [58,59] and Ho3þYb3þ or Yb3þTm3þ co-doped SiO2–LaF3 [60,61] glass-ceramics have been prepared through a similar sol-gel processing so far. The porosity of glass ceramics is an important factor for determining their optical properties. The precise control of the porosity can often enhance the interaction between light and materials. Here is an interesting attempt to increase photoluminescence intensity. SiO2–BaMgF4:Eu2þ glass-ceramic thin films were prepared by the sol-gel method using TFA [12]. Eu2þ was successfully doped in BaMgF4 nanoparticles dispersed in a silica glass matrix. The addition of DMF to a starting solution greatly influenced microstructure of the glass-ceramic films. That is, the film without DMF was dense while that with DMF was porous as shown in Figure 10.8. DMF is commonly used as a drying control chemical

(a)

(b)

1μm

Figure 10.8 FESEM images of the SiO2–BaMgF4:Eu2þ films (a) without and (b) with the DMF addition. Reproduced from reference 12 by permission of the Chemical Society of Japan

additive in preparing monolithic silica glasses in the sol-gel process. In contrast, DMF can be a cause of pore formation in the thin-film fabrication process where the coated solutions are put into the furnace kept at the heat-treatment temperature beforehand. Vaporization and/or decomposition of DMF from the inside of the film led finally to the porous microstructure. The films could exhibit blue emission peaking at 420 nm arising from the Eu2þ 5d ! 4f transition by the UV excitation at 290 nm. This emission was effectively enhanced in the porous film as shown in Figure 10.9, which was caused by multiple scattering of the excitation light.

10.5 10.5.1

Optical Properties Low Refractive Index and Anti-Reflection Effect

Alkaline earth fluorides having low refractive indices are useful, especially as coatings, in optics because of their high optical transparency in a wide range of wavelengths from IR to

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PL intensity (arb. unit)

700

λ ex = 290nm

600 500

without DMF

400

with DMF 300 200 100 0 350

400

450

500 550 600 Wavelength/nm

650

700

Figure 10.9 Photoluminescence spectra of the SiO2–BaMgF4:Eu2þ films without and with the DMF addition. Reproduced from reference 12 by permission of the Chemical Society of Japan

UV. Metal fluoride coatings that are prepared by common techniques such as evaporation or sputtering are normally dense and have refractive indices close to that of bulk fluoride materials. For a single-layer AR coating as shown in Figure 10.10, complete anti-reflection is attained when each optical parameter is under the following conditions: n21 ¼ n0 n2 n1 d ¼ l=4 where n0, n1 and n2 are the refractive index of the air, the coating and the component, respectively, d the thickness of the coating and l the wavelength of incident light [62]. Because the smallest refractive index is 1.40 (at 248 nm) of MgF2 among the nondeliquescent, practical inorganic materials, porous films are required to prepare single-layer AR coatings. Generally, the refractive index of materials decreases as the porosity increases according to the extended Lorentz-Lorenz formula [1]: 1 p ¼ ðn2 þ 2Þðn2m 1Þ= ðn2 1Þðn2m þ 2Þ Incident light (λ)

d

n0

air

n1

AR coating

n2

Figure 10.10

component

Schematic configuration of an AR coating

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where p is porosity and n and nm are the theoretical and the measured refractive indices, respectively. Optical design of a single-layer AR coating is therefore closely related to materials processing. Porous metal fluoride coatings are exclusively prepared by chemical solution deposition techniques including the sol-gel deposition method. The first work was done by Joosten et al. using a solution of magnesium trifluoroacetate or other fluorine-containing magnesium derivatives [63]. The solutions were deposited onto a fused silica substrate and then heated to 500–600 C. The resultant MgF2 films had excellent optical performance with relatively lower refractive index of 1.21. Thomas adopted colloidal suspension methods using methanol solutions of metal acetates or alkoxides and HF [64]. MgF2 and CaF2 suspensions were prepared as follows, MgðCH3 COOÞ2 þ 2HF ! MgF2 þ 2CH3 COOH MgðCH3 OÞ2 þ 2HF ! MgF2 þ 2CH3 OH CaðCH3 COOÞ2 þ 2HF ! CaF2 þ 2CH3 COOH For MgF2, the reactions were carried out in anhydrous methanol in which all the reagents were soluble and the product insoluble. In contrast, some water was required as cosolvent with the methanol to obtain complete solubility in the reaction of CaF2. Coatings were made on fused silica disks and polished calcium fluoride disks, which led to 45 % porous MgF2 and 50 % porous CaF2 films. Such porous coatings are also effective to enhance laser damage thresholds [65]. Larger colloidal particles (0.15 to 1 mm) of MgF2 are obtained from aqueous solutions of NaF and MgCl2 by aging at 80 C for 3 h [66]. Metal fluoride coatings can also be prepared through the trifluoroacetate-based sol-gel processing, as described above, using the dip- or spin-coating technique. In this case, their optical properties depend greatly on processing parameters. Figure 10.11 shows wavelength dependence of the refractive index of MgF2 films prepared using spin coating [67]. The coated films were dried at 80 C before annealing at 300, 400 or 500 C, which appeared to be effective to suppress abrupt vaporization of the solvents. It is seen that the higher heating temperature results in the lower refractive index. The porosity of the films heated at 300, 400 and 500 C were calculated to be 8.6, 28.7 and 39.4 %, respectively, using the extended Lorentz-Lorenz formula. It should be noted that the film thickness is relatively difficult to control because an increase in the number of coating layers in the deposition procedure may lead to inhomogeneity of the films in a depth direction. The heat-treatment temperature is not an only way to control optical properties, as mentioned above for the porous glass-ceramic thin films. The porosity and the optical thickness, n1d, of CaF2 films were controlled by adding organic compounds to coating solutions [68]. Organic additives such as aminoalcohols, 2-methoxyethanol, ethyleneglycol and cyclohexanol greatly influenced the deposition temperature and optical properties of porous CaF2 films because of the evaporation and/or decomposition of the additives during heating. Rywak and Burlitch prepared a nanocrystalline MgF2 sol, abrasion-resistant MgF2 films from it and SiO2–MgF2 composites both in the bulk and in the thin-film form [69]. In the synthesis of the composites, MgF2 sols were prepared from methanolic H2O2 and

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Mg(OCH3)2 treated in HF and mixed with silicate sols prepared from the acid catalysed hydrolysis and condensation of TMOS. Thin films of nanocrystalline MgF2 were prepared by spin-coating of the nanocrystalline MgF2 sol onto Si(100). The film thickness could be built up by repeating the coating procedure, but films thicker than 100 nm generally showed haziness caused by ‘islands’ on the surface. In contrast, using this multiple coating process, thicker SiO2–MgF2 composite films (ca. 150 nm) could be prepared on Si(100). These films had surface roughness of approximately –5 nm after being heated at 550 C. 1.60 300°C (d = 77 nm) 1.55

400°C (d = 70 nm) 500°C (d = 62 nm) Bulk MgF2

Refractive index

1.50

1.45

1.40

1.35

1.30

1.25

1.20

200

400

600

800

Wavelength/nm

Figure 10.11 Variation of the refractive index of the MgF2 thin films heat-treated at 300, 400 and 500 C for 10 min. Reproduced from reference 67 by permission of Springer

Murata et al. used MgF2 sols obtained from magnesium acetate and HF to prepare ultralow refractive index MgF2 thin films [70–72]. The sols were autoclaved in a Teflon cell at 100–180 C and then coated on SiO2 glass or CaF2 crystal substrates by spin-coating. Subsequently, the coated films were heat treated at 150 C. As shown in Figure 10.12, the MgF2 films exhibit the AR effect over a very wide range of wavelengths in the VUV range. They also fabricated porous MgF2–SiO2 thin films consisting of MgF2 particles connected by an amorphous SiO2 binder [73]. The films had a low refractive index of 1.26, sufficient strength to withstand wiping by a cloth and a high environmental resistance. The films could be uniformly formed on curved substrates and at relatively low temperatures, such as

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100 C. Because the wet-processed fluoride coatings generally have lower mechanical strength than the dry-processed ones, the nanocomposite with SiO2 may be the best way to produce AR coatings with improved performance. Sermon and Badheka reported the preparation of MgF2 xerogels with high surface area (821 m2/g at 423 K), high pore volume (1.00 cm3/g) and small particle size (5–7 nm) [74]. They could be readily applied to substrates by traditional methods to give high porosity xerogel coatings. Recently, Kru¨ger et al. reported the deposition of MgF2 films of optical quality through an anhydrous low temperature sol-gel synthesis using MgF2 sols [75, 76]. The MgF2 precursor solution was prepared from a suspension of magnesium methoxide in methanol and a nonaqueous HF solution in methanol. The solution was then deposited onto Si substrates by spin coating. The refractive index of the resultant films was smaller than those of the bulk phase due to a lower density of the MgF2 films as compared to the bulk material. Quite an interesting approach has been made recently by Grosso et al. to prepare magnesium oxyfluoride-based ultralow-dielectric-constant optical thin films [77]. Precursor solutions were prepared simply by mixing magnesium acetate, water and TFA. Highly porous and resistant semicrystalline magnesium oxyfluoride thin films were obtained through liquid deposition followed by a flash and short thermal treatment. The TEM observation revealed that the film was composed of rigid and atypical porous inorganic networks made of coalesced hollow particles and vesicles. Lately, MgF2-based optical thin films exhibiting refractive indices ranging from 1.08 to 1.2 were reported [78]. The generation of the porosity (Figure 10.13) responsible for the ultralow refractive indices of these materials was thought to be triggered by the thermal decomposition of metallic precursor ligands. 100.0 99.0 98.0 Transmittance / %

97.0 96.0 95.0 94.0 93.0 92.0 91.0 90.0 140.00

150.00

160.00

170.00

180.00

190.00

Wavelength / nm

Figure 10.12 Transmittance of CaF2 substrate with antireflection coatings made from the autoclaved MgF2 sol on both sides. Reproduced from reference 70 by permission of Springer

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Figure 10.13 TEM images of a typical vesicle-like structure characteristic of the material with 3D magnesium oxofluoro network. Reproduced from reference 78 by permission of the American Chemical Society

10.5.2

Luminescence

Luminescence of rare-earth ions doped into inorganic solid materials has found numerous practical applications in phosphors and lasers [79]. Doped rare-earth ions can be excited by electromagnetic radiations, cathode rays, electric voltages or any other energy sources through intraconfigurational f–f transitions of inner 4f electrons, interconfigurational f–d transitions, charge transfers between rare-earth ions and anions or host lattice absorptions. Luminescent behaviour of excited rare-earth ions is subject to crystal fields of host lattices and also to lattice imperfection. Furthermore, lattice phonon affects quantum efficiencies of emissions from excited rare-earth ions. It is interesting and important to design, synthesize and utilize host lattices containing fluorine for desired luminescence because fluoride crystals generally give lower phonon energies than oxide crystals. Trivalent europium (Eu3þ) ions show red luminescence due to intraconfigurational f–f transitions. A comparison between LaF3 and LaOF makes it easy to understand effects of host crystal lattices on luminescence. In the excitation, as well as intraconfigurational f–f excitations of inner 4f6 electrons, interconfigurational charge-transfer excitations, 4f6 ! 4f7L1 (L is ligand), can take place in the Eu3þ ion. The charge-transfer excitations are allowed, are usually positioned in the UV region and have much higher intensity than the f–f excitations [79]. Figure 10.14 shows excitation spectra for a 5D0 ! 7F2 (611 nm) emission of La0.9Eu0.1F3 and La0.9Eu0.1OF thin films prepared by the sol-gel method using TFA [10,80]. La0.9Eu0.1OF exhibits a strong excitation band at 272 nm, resulting from an O2–Eu3þ charge transfer [81]. On the other hand, the charge-transfer excitation is not

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Emission: 611nm

Intensity (arb. unit)

(5D0→7F2)

La0.9Eu0.1OF

La0.9Eu0.1F3

240

280

320 360 400 Wavelength / nm

440

480

Figure 10.14 Excitation spectra for a 5D0 ! 7F2 (611 nm) emission of Eu3þ in La0.9Eu0.1F3 and La0.9Eu0.1OF thin films. Reproduced from reference 10 by permission of the American Ceramic Society

observed in La0.9Eu0.1F3 in the measured wavelength regime. The position of the chargetransfer band generally shifts to higher energy for decreasing covalency in chemical bonds. It takes much more energy to remove an electron from an F ion in La0.9Eu0.1F3 than from the O2 ion in La0.9Eu0.1OF. Lanthanum oxyhalides including LaOF doped with luminescent rare-earth ions were studied as phosphors for potential applications in the earlier times [82]. Crystal structures and optical or luminescent properties of rare-earth oxyfluoride compounds were extensively investigated in the 1990s [81–87]. Usually, rare-earth doped oxyfluoride materials are synthesized by heat treatments of rare-earth fluorides in the air [88] or mixtures of rareearth oxides and NH4F in N2 atmosphere [89]. Eu3þ-doped YOF and LaOF hollow spheres were synthesized by a facile template route [90]. The templating carbonaceous spheres were prepared and dispersed in metal nitrate solutions. An NH4F solution was then dropped in. The resultant suspension was aged and heated to 700 C. Emissions from Eu3þ were typical in rare-earth oxyfluoride structure. Du et al. prepared monodisperse doped lanthanide oxyfluoride (Tb3þ, Ce3þ, Eu3þ, Yb3þ or Er3þ-doped LaOF and GdOF) nanocrystals by codecomposing lanthanide trifluoroacetate (Ln(CF3COO)3) precursors in oleic acid/oleylamine [91]. The nanocrystals showed controlled sizes (2–7 nm) and shapes (nanopolyhedra and elongated nanocrystals) and could form a large-area superlattice on copper grids via self-assembly. Luminescent behaviours of differently sized and shaped nanocrystals were revealed to be largely dependent upon their microstructures due to small size effects and surface-structure effects. Armelao et al. developed fluorinated -dichetonate compounds La(hfa)3diglyme (Hhfa ¼ 1,1,1,5,5,5hexafluoro-2,4-pentanedione; diglyme ¼ bis(2-metoxyethyl)ether) and employed them as

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single-source precursors in the sol-gel processing of lanthanum fluoride and oxyfluoride thin films [92]. As shown in Figure 10.15, LaF3 and LaOF thin films prepared were highly transparent in the 200–2500 nm wavelength range and expected to be utilized in the realization of visible and infrared emitting materials and devices by doping with a variety of rare-earth ions.

90 Transmittance (%)

300°C 500°C 700°C

80

900°C 70

60

50 400

800

1200 1600 2000 Wavelength (nm)

2400

Figure 10.15 Transmittance spectra in the 200–2500 nm wavelength range for the lanthanum oxyfluoride films annealed for 1 h in air at different temperatures. Reproduced from reference 92 by permission of the American Chemical Society

These days, the development of multiphoton phosphors that employ novel conversion schemes to provide efficient visible emissions is highly required using, for example, quantum cutting phenomenon. Then the sol-gel synthesis was explored to obtain a LiGdF4:Eu3þ phosphor which could have a quantum efficiency of 190 % [36]. The samples prepared in the system LiGd1xEuxF4 showed the characteristic Eu3þ emission upon UV irradiation in one of the absorbing level of Gd3þ (6I7/2, 6P5/2 or 6P7/2) or upon blue light excitation in the 5D2 manifold of the Eu3þ ions. Lately, it has been demonstrated that sol-gel-derived LiGdF4:Eu3þ powders can show emission of two visible photons per absorbed VUV photon [93]. This mechanism is explained by a two-step energy transfer when exciting Gd3þ ions in their 6GJ high energy level. Divalent europium (Eu2þ) ions can show 5d ! 4f emissions in a wide wavelength range because energy levels of outer 5d electrons are largely affected by the crystal field strength. Eu2þ-activated BaMgF4 thin films were easily obtained by the sol-gel method from trifluoroacetate solutions containing Ba2þ, Mg2þ and Eu3þ [41]. Heat treatments of the films at temperatures of approximately 650 C in flowing nitrogen atmosphere led to the reduction of Eu3þ to Eu2þ, which was accommodated in the BaMgF4 structure. It is somewhat difficult to synthesize rare-earth-doped complex fluorides because the decomposition temperature of each metal trifluoroacetate is slightly different. Thus the intensity of the blue-violet luminescence of Eu2þ depended sensitively on the heating temperature,

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heating time and the europium concentration. The strongest emission was observed with the BaMgF4:Eu2þ (15 mol %) film that was heated at 650 C for 10 min. Perovskite-type BaLiF3:Eu2þ was synthesized in a similar way [42]. The trivalent Eu3þ, which was used as the acetate for the starting material, was reduced to divalent Eu2þ in the pyrolysis process of BaLiF3. Figure 10.16 shows photoluminescence spectra of the BaLiF3:Eu (3 at %) film. The excitation wavelength used was 254 nm. A broad blue emission with a peak wavelength of 408 nm is due to the 4f65d ! 4f7 transition of excited Eu2þ. The concentration quenching of the blue emission occurred at 5 % of Eu2þ in BaLiF3, indicating that Eu2þ was homogeneously dispersed in the BaLiF3 host lattice. The Eu3þ ! Eu2þ reduction is generally achieved by firing materials at high temperatures in a reducing atmosphere or air. When H2 gas is used, a reduction easily proceeds as follows: 2Eu3 þ þ H2 ! 2Eu2 þ þ 2Hþ In contrast, the reduction in the air should be caused by chemical interactions between Eu3þ and constituents of host lattices. Peng et al. explained the reduction of Eu in Sr4Al14O25 based on a charge compensation model where strontium vacancies played a key role [94]. The Eu reduction seems to be more feasible in fluoride lattices due to chemical interactions between Eu3þ and F. When Eu3þ is doped in the BaLiF3 lattice through a solid-state reaction between BaF2, LiF and EuF3, it is reduced to Eu2þ by a following reaction, which was actually achieved at 750 C in the flowing nitrogen atmosphere [95],

Intensity (arb. unit)

Eu3 þ þ F ! Eu2 þ þ F

350

400

450

500

550

600

650

700

Wavelength / nm

Figure 10.16 Photoluminescence spectra of the BaLiF3:Eu2þ (3 at%) film. The excitation wavelength used was 254 nm. Reproduced from reference 42 by permission of Elsevier

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The decomposition reaction of the metal trifluoroacetates involves the generation of gaseous phases such as (CF3CO)2O, CF3COF and COF2 [19] in addition to the dehydration and the vaporization of the organic components. It seems that F can be supplied from these organofluorine species because, in the presence of water, F ions can be supplied by hydrolysis of CF3COF and COF2. Thus the Eu3þ ! Eu2þ reduction process is regarded as a solid–gas reaction in this case. Eu-doped CaF2 and MgF2 films were synthesized by the sol-gel method using TFA [96]. It was found that the films of CaF2 and MgF2 doped with Eu resulted in luminescence of Eu2þ and Eu3þ ions, respectively. Eu2þ-activated CaF2 and SrF2 nanoparticles were also prepared by the sol-gel method assisted with a thermal-carbon reducing atmosphere (TCRA) treatment [97]. Photoluminescence with a peak at 425 and 416 nm due to Eu2þ was observed for CaF2 and SrF2, respectively. Li et al. reported systematic manipulation of the morphologies and architectures of NaYF4 microcrystals using a simple and mild solution-growth method [98]. They investigated influences of fluoride sources (NaF and NH4F) and pH values on the shapes of NaYF4 microstructures developed during hydrothermal treatments. In addition, experimental results indicated that optical properties of -NaYF4:Tb3þ phosphors with different microarchitectures were strongly dependent on their morphologies and sizes. Nano-sizing and microstructure engineering are now becoming key to the fabrication of novel luminescent materials both in oxides and non-oxides.

10.6

Concluding Remarks

This chapter reviewed recent advances in sol-gel technologies for the elaboration of novel, structure-controlled and high-performance metal fluoride nanomaterials in the form of nanocrystals, thin films, nanocomposites and oxyfluoride glass ceramics. Their excellent optical properties, which can surpass those of metal oxides, mean that metal fluorides are increasingly studied by chemists, physicists and materials scientists. The sol-gel method described in this chapter includes the direct fluorination of oxide precursor gels, the thermal decomposition of metal organofluorine materials and the building-up of metal fluoride crystals from F ions and metal cations in the solutions. Although the functions mentioned in this chapter were limited to antireflective effects and luminescence, there is no doubt that these fluoride nanomaterials have a great potential for use in optics, photonics and optoelectronics. Energy and environmental materials are also emerging from fluoride nanomaterials, which can also be fabricated by the sol-gel method.

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11 Fluoride Glasses and Planar Optical Waveguides Brigitte Boulard Laboratoire des Oxydes et Fluorures, UMR CNRS 6010, Universite´ du Maine, avenue O. Messiaen, 72085 Le Mans, cedex 9, France

Fluoride glasses are of great interest for optical applications, as low phonon energy hosts for rare-earth ions. With an increasing need for more compact optical devices that impose higher dopant concentration, research activity has been focused on the fabrication of fluoride glass planar waveguides to produce integrated lasers and optical amplifiers, especially those based on Er3þ-doped glasses. The aim of this review is to give the state of the art of fabrication technology for fluoride glasses and to present the most interesting results obtained on confined waveguides. Recent developments in fluoride transparent glass ceramic waveguides will also be mentioned.

11.1

Introduction

Fluoride glasses have been known since 1926, when it was discovered that beryllium fluoride could be cooled below the liquidus temperature without crystallization. Little further development took place until 1974, when Poulain et al. made the first synthesis of ZrF4-NaF-BaF2-NdF3 [1]. During the past three decades, much attention has been paid to the family of fluoride glasses [2], especially fluorozirconates such as ZBLAN

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Ó 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

332

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glass (ZrF4-BaF2-LaF3-AlF3-NaF), because of their potential in the development of optical fibres for long-distance telecommunication links. These applications require the absence of optical attenuation from scattering by crystals and particulates, but also ultrahigh purity materials containing no light-absorbing impurities such as transition metal (Fe2þ, Cu2þ, etc) hydroxyl (OH) ions. To date, the ultimate predicted ultralow optical loss of 0.01 dB/km compared to 0.2 dB/km for silica glasses [3] has not been reached, the biggest obstacle being in reducing transition-metal impurities and inhibiting formation of crystallites during the melt cooling operation or after reheating above the glass transition temperature (Tg). Nevertheless, fluoride glass remains an attractive material in shorter optical devices with applications lying in the visible and mid-IR spectral range, including lasers and amplifier operating at wavelengths not accessible with silica-based glasses, thank to their low phonon energy (500–600 cm1) compared to silica (1100 cm1). In particular, the development of optical communication necessitates the design and manufacture of integrated optic lasers and amplifiers, especially of those based on erbium-doped glasses. Integrated optical amplifiers (IOA) are used to bring the fibre to the home because they gather two complementary criteria: the high debit rate by use of wavelength distribution multiplexing (WDM) system and local network architecture. Indeed they inherit all the advantages of erbium doped fibre amplifiers (EDFA) – weak noise, weak polarization effect, absence of interference between channels in WDM application, in contrast with semiconductor optical amplifiers (SOA) [4]. The short length of integrated amplifiers imposes higher Er3þ concentration and higher pump-power density than fibre so the choice of the glass matrix is particularly critical. Actually, the rare-earth solubility is strongly related to the crystal chemistry of the glass. Except BeF2, fluoride glasses offer high coordination sites for rare earth ions and thus represent a unique optical host for rare earth ions; solubility can reach 10 mol%, depending on the ion size [5] while solubility is only a few hundred ppm in the tetrahedra-based network of silica. In this chapter, the current status of the processing technologies in the planar waveguides fabrication and their performance as integrated optical amplifier and laser is reviewed. A special part will be dedicated to waveguides based on transparent glass ceramics, an emerging material in the field of active optics, as it may offer macroscopic glass properties and crystal-like spectroscopic characteristics [6].

11.2

Rare Earth in Fluoride Glasses

The composition and thermal and physical properties (namely, Tg, stability criteria DT, refractive index and vibration frequency of metal-F bond) of some fluoride glasses that have been used for waveguide fabrication are given in Table 11.1. One can note the relatively high concentration of rare-earth fluorides, at least 5 mol%. (with YF3 assimilated to the lanthanide family REF3). Such a high concentration allows more flexibility in rare-earth combinations for active optical applications, especially for upconversion processes and enhanced pump absorption due to energy transfer between lanthanide ions.

Fluoride Glasses and Planar Optical Waveguides

333

Table 11.1 Chemical composition, thermal and physical properties of fluoride glasses used to achieve planar waveguides. o is the vibration frequency that accounts for the multiphonon absorption acronym ZBLAN ZBLANPb ZBLA ZELAG BIG BIGNa PZG IZSB ALF70

11.2.1

Composition (mol%)

Tg (°C)

DT (°C)

nD

o (cm1)

52ZrF4-24BaF2-5LaF3 4AlF3-15NaF 52ZrF4-19BaF2-5PbF2 5LaF3-4AlF3-15NaF 57ZrF4-34BaF2-5LaF3 4AlF3 60ZrF4-26LaF3-6ErF3 1AlF3-6GaF3 25BaF2-23InF3-12GaF3 25ZnF2-15GdF3 15BaF2-18InF3-12GaF3-20ZnF2 10YbF3-10GdF3-15NaF 36PbF2-24ZnF2-35GaF3 5YF3-2AlF3 16BaF2-40InF3-20ZnF2 20SrF2-4GdF3-2NaF-6GaF3 37AlF3-15CaF2-12MgF2-6BaF2 15YF3-9SrF2-6NaPO3

272

90

1.504

580

254

81

1.517

580

307

85

1.516

580

392

49

1.503

600

324

145

1.505

510

310

140

1.487

510

270

52

1.577

560

300

90

1.503

510

435

130

1.432

<650

Fundamentals

Rare earths (REs) are characterized by [Xe] 4fn 6s2 electronic configuration and the most stable oxidation state is þIII with the 5s and 5p electrons remaining untouched. These electrons act as a screen for 4f electrons toward the surrounding environment, so that the energy levels of the ion are almost independent of the host. Unlike the energy levels, the transition probabilities between 4f states are host sensitive through phonon energy and clustering due to concentration effect. The quantum efficiency of a given transition is limited by nonradiative relaxations that may be classified as (i) multiphonon relaxations, (ii) cross relaxations and (iii) upconversion processes. The probability of multiphonon relaxations WMP is a function of energy gap DE to the next lower level [7]: WMP ¼ B exp ðDEÞ with B and as constants that depend only on the host. It is commonly admitted that WMP is negligible when DE > 7o. As an example, the lack of emission at 1.3 mm for Pr3þ in silica -based glasses is a direct consequence of this, the energy gap being 3000 cm1. Figure 11.1 shows the energy levels of rare-earths and emissions of primary interest with regard to visible, NIR and mid-IR applications in fluoride glasses.

Functionalized Inorganic Fluorides

334

2

P11/2

30 1 4

G11/2

P0

1

4

4

I9/2

H4 3+

Pr

Nd

3+

3

0,80

2

F7/2

0,85

H6

Er

F5/2

F4

3

3+

2

H5

3

I15/2

1,20 1,47 2,25

1,06

1,34

I13/2 4 I11/2

3

2,75

I13/2

H4

1.03

H5

4

I15/2

4

3

1,88

3

4

I11/2

0,41

H6

4

0,99

F3/2 0,93

F4

3

F9/2 I9/2

4

1,31

G4

4

4

1

3

0

H11/2

0,77

0,89

0,63

2

1,55

D2

0,65

H11/2 4 S3/2

1

10

G4

2

0,48

3

H9/2

P3/2

20

1,08

Energie (103.cm)–1

2

D2

2

3+

Tm

Yb3+

Figure 11.1 Energy level diagram of RE3þ ions (RE ¼ Pr, Er, Tm, Nd, Yb) of main interest for visible and IR applications. Yb3þ is used as a sensitizer because of its high absorption cross section and the possible energy transfer to its neighbouring ions

11.2.2

Applications: Laser and Optical Amplifiers

In all active optics applications, rare-earth doped fibres or waveguides possess significant advantages over bulk solid-state lasers because of long interaction length in laser media [8]. Laser action in waveguiding structures occurs at a threshold well below the bulk one, thus allowing higher gain and consequently large population inversion with low pumping power. The properties of planar optical amplifiers and lasers are expected to be similar to those of the fibre because of confinement and moreover offer the possibility of integrating other optical functions. A demonstration was made for the first time in 1974 with a Nd3þ-doped integrated optical glass laser, while the same material (a borosilicate glass) was employed as an early fibre laser [9, 10]. Some activity on RE- doped fluoride glass is telecommunication oriented with research devoted to broad-band optical amplification, within the loss-low window (1.2–1.6 mm) of silica optical fibre. The attenuation curve of highly transparent silica fibre (termed ‘allwave’) shown in Figure 11.2 reveals Pr3þ, Tm3þ and Er3þ to be of great interest. Erbium-doped amplifiers (EDFA) based on silica can cover the 1.5–1.6 mm band but signal amplification in the entire 1.2–1.6 mm window cannot be achieved; fluoride glass host is required at 1.3 mm with Pr3þ and 1.48 mm with Tm3þ. Moreover, fluoride erbium-doped fibres are of practical interest because they possess a flat and broadband gain bandwidth as a function of signal wavelength [8]. This property is fundamental for wavelength division multiplexing (WDM), which allows the transmission of several wavelengths simultaneously over a wide range of energy. It is expected that WDM transmission will also be present in local networks, with integrated optical devices allowing complementary functions such as amplification, division and multiplexing.

Fluoride Glasses and Planar Optical Waveguides

335

ATTENUATION (dB/km)

2.0 1,31 µm (Pr3+)

1.0 0.8

(OH) 1,47 µm (Tm3+)

1,55 µm (Er3+)

0.6 0.4 Allwave 0.2 1.0

1.1

1.2 1.3 1.4 1.5 WAVELENGTH (µm)

1.6

1.7

Figure 11.2 Attenuation of silica optical fiber and operation wavelengths for optical amplification. ‘‘Allwave’’ refer to OH free fiber developed by Lucent Technology

Besides telecom applications, RE-doped fluoride glass waveguides can also be used as planar lasers in the area of high range power diode-pumped solid-state lasers. The idea is to improve the beam quality by delivering a diffraction-limited beam from a single-mode waveguide. Lasings have been obtained for a large number of rare-earth transitions in fluoride glass fibres from blue to the mid-IR (see Table 11.2), including several three-level systems that hardly lase in bulk configuration. As already noted, fluoride glass fibre allows

Table 11.2 efficient emissions of rare-earth ions RE3þ observed in fluoride glass fiber lasers or amplifiers (from [6,13]). DE is energy gap to the next lower level from the emitting level. Yb3þ Codoping (cod.) is used for pumping at 0.98 mm RE3þ Pr

Tm Er

Ho Nd

l (mm) 1.3 (IR) 0.605–0.635 (red) 0.52 (green) 0.49 (blue) 1.48 (NIR) 0.45 (blue) 3.45 (mid-IR) 2.79 (mid-IR) 1.54 (NIR) 0.65 (red) 0.545 (green) 3.9 (mid-IR) 0.55 (green) 1.34 (NIR) 1.06 (NIR)

transition

lexc (mm)

G4 ! 3H5 P0 ! 3H6 3F2 3 P1, 1I6 ! 3H5 3 P0 ! 3H4 3 H4 ! 3F4 1 G4 ! 3H6 4 F9/2 ! 4I9/2 4 I11/2 ! 4I13/2 4 I13/2 ! 4I15/2 4 F9/2 ! 4I15/2 4 S3/2 ! 4I15/2 5 I5 ! 5I6 5 S2, 5F4 ! 5I8 4 F3/2 ! 4I13/2 4 F3/2 ! 4I11/2

1.02 0.48, 0.85apte (Yb cod.) 0.48, 0.85apte (Yb cod.) (1.02 þ 0.83)uc 1.4 1.12, 0.98apte (Yb cod.), 0.65uc 0.65 0.98 1.48, 0.98apte (Yb cod.) 1.48uc 1.48, 0.98et (Yb cod.), 0.80uc 0.65 0.88uc 0.80 0.80

1 3

uc: up-conversion pumping. apte: addition of photon by energy transfer pumping.

DE (cm1) 2900 3800 4400 3800 4100 6100 no no no 2200 2800 no 3000 5500 5500

336

Functionalized Inorganic Fluorides

emission at wavelengths not achievable with oxide glasses [11]; for example mid-IR emissions located in the 3–5 mm atmospheric window are specific to fluoride glasses doped with Er3þ at 2.8 mm and Ho3þ ions at 3.9 mm, taking advantage of fluoride glasses’ transparency at wavelengths greater than 5 mm while silica transmission starts decreasing at 3 mm. Because of the strong confinement of the light in fibre or channel configuration, upconversion processes can be nicely optimized to achieve powerful and compact new visible laser sources, i.e. blue (with Tm3þ at 453 nm), green (with Er3þ or Ho3þ at 550 nm) and red (with Er3þ at 650 nm). A low phonon energy material is generally required to give access to most of the involved metastable excited states. This is why only RE-doped fluoride materials have been successfully used. Some applications require the simultaneous generation of the three primary light colours (red, green and blue) allowing additive synthesis of light in the visible spectrum [12], to be used for miniature video-projection; in this case, Pr3þ or Tm3þ,Er3þ codoped fluoride glasses appear attractive. Table 11.2 gathers interesting wavelengths for laser action and optical amplification in visible and IR observed in fluoride glasses. An exhaustive list of results can be found in [13].

11.3

Fabrication of Waveguides: A Review

There is no single waveguide fabrication process that is universally applicable to the whole range of glass materials. Various techniques are available and some have been used to produce fluoride glass planar waveguiding structures: ionic exchange, chemical and physical vapour deposition (CVD and PVD), pulsed laser deposition (PLD), radio frequency (RF) magnetron sputtering and electron-beam sputtering (EB). The main difficulties encountered in these processes are (i) to reproduce the composition and optical properties of multicomponent fluoride glass; (ii) to avoid crystallization during the quenching step or after reheating the glass (over Tg) needed for ion-diffusion and sintering, (iii) to get low propagation losses. Moreover the film thickness has to be achievable in a reasonable processing time for a mass production. The main advantages and drawbacks of the processes mentioned above as applied to fluoride glasses are listed below: • Ionic exchange. The F ion is replaced by Cl, OD coming from gaseous phase [14,15] or alkali cations (Liþ, Kþ, Agþ) are exchanged from a molten salt [16,17] or by dry interdiffusion [18]. The maximum refractive index gradient achieved by anionic and cationic diffusion on fluoride glass is given in Table 11.2. The optical properties of the starting glass may be affected by the change in size of the exchanged ions by generating stress at the surface. • CVD. After its success in the preparation of low-loss silica based fibres, this process has been adapted to fluoride materials using volatile fluoride containing organometallic precursors ( diketonates) and HF or SF6 fluorinating gases to prepare ZrF4-based fluoride glasses [19]. Direct synthesis of transparent thin film on a substrate has been achieved by plasma-enhanced CVD [20]. The main disadvantage of using precursor organic ligands is the presence of residues containing carbon and/or oxygen that cause quenching of luminescence and defect centres.

Fluoride Glasses and Planar Optical Waveguides

337

• PVD. The noncongruent evaporation of multicomponent glass restrains the choice in fluoride glass composition [21, 22]. As for all deposition techniques, the thermal expansion coefficients of the substrate and of the deposit have to be close to avoiding film cracking. The technique allows high growth rate (0.1 mm/min). • Sol gel. The synthesis usually involves wet chemistry reactions and is based on the inorganic polymerization of molecular precursors – either organic like alkoxides or inorganic like nitrate – for active optical application. Thin films are produced directly from the solution by dip or spin coating and can be converted in a fluoride glass film by fluorination with HF at temperature below the Tg (200 °C) [23, 24]. The critical step is the drying to remove OH groups. • Spin casting. The low viscosity of fluoride glass at its melting point allows to use this technique, provided that the glass is rapidly quenched to avoid crystallization of the molten glass layer [25]. Fast quenching enhances thermal stress at the film-substrate interface. To overcome this difficulty, the substrate is heated close to the glass transition temperature. The main drawback of the technique is the great dependence of film thickness and uniformity with spin speed parameters. • Sputtering. Arþ plasma or electron beam are used to sputter atoms from the target to the subtsrate. Fluoride thin film coatings are generally difficult to deposit by sputtering processes because deposition rates are low and coating quality is also generally low compared to films deposited by evaporation [26]. The evaporation of the glass has to be congruent for electron beam sputtering [27] unlike RF sputtering. Processing gas (SF6,. . .) is introduced during deposition in order to compensate fluorine that is frequently deficient at the surface target [28]. • PLD. This is a versatile materials fabrication technique that enables to preserve the stoichiometry of multi-component targets. The films are grown by ablation of fluoride glass, resulting in a congruent vaporization. Fluorine F2 is used as processing gas [29]. A basic glass waveguide consists of a guiding core glass surrounded by a smaller refractive ‘cladding’ material (fluoride glass or single crystal like CaF2. . .). The refractive index difference Dn between the core and the clad is an important parameter because it governs the guiding properties of the waveguide. Often, the top surface of the waveguide is left open to the air to make the fabrication simpler. In order to minimize coupling loss between the optical fibre and the waveguide, single-mode structure is preferred. for which the maximum waveguide thickness e is given by: ﬃ e ep qﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ n2g n2s ¼ l l

qﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ p Dnð2ng DnÞ 2

with ng and ns the guide and substrate refractive indices. Flexibility in Dn is achieved by using core and clad glasses with close compositions. Despite a refractive index that often does not match the required index, single crystal is chosen because it provides injection faces of optical quality by cleavage. Table 11.3 compares the optical properties of different fluoride glass waveguides. Although being the most promising materials owing to their low phonon energy, production of InF3-based channel waveguides with acceptable loss has been unsuccessful due to

338

Functionalized Inorganic Fluorides

the low Tg compared to the temperature required for the diffusion mechanism [16]. The process based on spin casting of fluoroaluminate glass currently holds the record for the minimum propagation losses, of less than 0.1 dB/cm. Table 11.3 Comparison of fabrication methods and characteristics (refractive index change, thickness and propagation loss) of fluoride glass planar waveguides Technique

Waveguide glass

Substrate

PVD

PZG ZELA 60ZrF4-35BaF2-5PrF3 79ZrF4-8AlF3-13NaF(b) ZBLA ALF70 ZBLANPb ZBLAN IZSB

CaF2(a) CaF2 CaF2 BaF2(a) CaF2 ALF101(c) ZBLAN MgF2(a) exchanged ions Agþ/Naþ

BIGNa

Liþ/Naþ

BIG

OD/F

ZBLA

Cl/F

CVD EB sol gel spin casting PLD

ionic exchange

Dn core/clad 0.14 0.10 0.09 – 0.08 0.21 0.02 0.10 Max. Dn gradient 0.03 2 h–300 °C 0.19 16 h–300 °C 0.07 100 h–347 °C 0.07 10.5h–250 °C

e (mm)

Loss Ref. (dB/cm)

1–10 0.4 [32] 1–5 1.2 [22] 4–7 15.9 [36] <5 – [28] 25 – [24] 2–10 0.1 [20] 50 – [33] 1 5.1 [30] depth loss Ref. (mm) (dB/cm) 20 2.4 [18] 5

very high

[17]

22

5

[16]

8

0.3

[14]

(a)

refractive index of single crystal nD: CaF2 ¼ 1.434; BaF2 ¼ 1.474; MgF2 ¼ 1.378. ZBLAN target. (c) modified ALF70 composition with low refractive index: nD ¼ 1.417. (b)

11.4

Performance of Active Waveguides

The successful fabrication of planar waveguides with proper spectroscopic properties and low propagation loss is not the ultimate goal. Demonstration of efficient laser or optical amplification requires a channel waveguide configuration, with lateral and transverse confinement. This supplementary step often generates extra propagation loss. Figure 11.3 shows the possible geometries of confined structures. Except in the case of laser writing and micromachining, processing using photolithography is necessary to achieve these geometries. Photolithography, which is well established for silica, is not directly transferable to fluoride material because of (i) incompatibility between the glass and the solvent used to develop the resin; and (ii) thermal resistance and chemical inertia of the mask toward the glass components and sometimes reactive atmosphere. Obviously, this is one main reason why the literature on the production of fluoride glass-channel waveguides is not as extensive as that for planar waveguides.

Fluoride Glasses and Planar Optical Waveguides

339

resist SiO2 bulk glass

SiO2 after RIE

ionic exchange focused laser beam

resist removed

substrate

mask removed

bulk channel WG

focused laser beam

channel WG

channel WG

buried channel WG

(a)

(b)

resist planar WG resist

resist SiO2 substrate

substrate RIE or wet etching film deposition resist

substrate

SiO2 after RIE

film deposition resist removed

resist removed strip WG

ridge WG

charged WG

(c)

(d)

(e)

Figure 11.3 Different geometries of confined waveguides (WG): (a) laser written WG (direct writing or self writing), (b) exchanged WG (buried or not buried), (c) strip WG (‘‘lift-off’’ process), (d) ridge WG and (e) charged WG. The resist mask in (b), (c), (d) and (e) is obtained by standard photolithography. RIE is reactive ion etching

For F/Cl exchange, the use of a silica mask leads to the best results in terms of optical quality of the guide [30] owing to its chemical inertia against the glass and the reactive atmosphere, in comparison to a metallic mask; final burial by HF treatment results in an elliptical waveguide. In deposition techniques, a resin mask is acceptable when deposition temperature is low (i.e. < 200 °C); as an example, the ‘lift-off’ process have been successfully applied to get PZG strip waveguides by PVD [31]. For ridge waveguide fabrication, wet (chemical) or physical etching are available techniques. Wet etching of fluoride glass requires specific solutions, i.e. ZrOCl2/HCl for ZrF4-based glass [32, 33], AlCl3/HCl for AlF3-based glass, which prevent precipitation by forming a stable fluoride complex. The

340

Functionalized Inorganic Fluorides

main problems encountered with wet etching are the nonrectangular profile and the extensive roughness of the edges [34]. Micromachining has been applied to generate ridge-like waveguides between two parallel grooves structured by laser ablation [29], also leading to poor edge quality. One alternative may be to use reactive ion etching (RIE), but a test performed on ZrF4-based glass was not conclusive [35]. On the whole, ionic exchange is considered to be the most suitable low-cost technique to fabricate channel waveguides. To complete the overview, laser writing of channel waveguides has to be mentioned. This is a very fast and powerful technique that can be applied both to films and bulk materials. Irradiation of an area in the glass is used to achieve photo-induced refractive index changes, up to 0.01. Practically, waveguides are written in glass with focused fentosecond laser pulses (i.e. l ¼ 800 nm [36]), or continuous waves in the case of films to avoid damage (i.e. l ¼ 244 nm [37]). Another possibility is to create a confined waveguide between two photothermal expansion ridges (i.e. l ¼ 244 nm [38]); this way, the guiding zone is sheltered from possible radiation damage that can strongly affect the guiding properties.

11.4.1

Optical Amplifier

The most important characteristic to indicate the performance of an optical amplifier is the net or external gain that takes into account the coupling loss (C), propagation loss () and RE-absorption a. It is given by the following equation: Gnet ðdBÞ ¼ Gon=off 2C ðL þ a NLÞ where L is the waveguide length and N the RE concentration. The relative gain Gon/off, i.e. the ratio of the output signal when the laser pump is on and off, is a quicker measurement and thus gives a preliminary estimation of amplifier performance (Figure 11.4). 4

Gon/off (dB/cm)

3

2

1

0

ZBLA PZG 0

50

100

150

200

250

P @ 980 nm (mW)

Figure 11.4 ‘On/off’ gain at 1.53 mm as function of the incident pump power in Er3þ-doped ZBLA and PZG fluoride glass waveguides pumped at 980 nm. Inserted in the graph, near field images showing single mode (ZBLA) and multimode (PZG) waveguides at this wavelength

Fluoride Glasses and Planar Optical Waveguides

341

Table 11.4 summarizes the most interesting results for fluoride glass integrated amplifiers. As can be noted, both PVD and ionic exchange technologies, developed simultaneously by two French teams in Rennes and Le Mans, have shown an ability to make waveguide amplifiers. Although waveguides produced by PVD exhibit similar or even higher relative gain, net gain has been observed only for Cl/F exchanged ZBLA in a single-mode channel waveguide (Gnet ¼ 2.5 dB for L ¼ 1.9 cm with 1.48 mm pumping). This can be partially explained by the mismatch between the optical modes of the coupling fibre and the smaller mode of the PZG ridge waveguide. Table 11.4 characteristics and performance of fluoride glass waveguide amplifiers at 1.53 and 1.05 mm. The lifetime written in italic corresponds to the value measured in bulk l (mm)

Glass (RE mol%)

WG section (form)

Gon/off (dB/cm)

Ppump (mW)

t (ms)

loss (dB/cm)

Ref.

1.53

PZG 1 Er3þ-1Ce3þ ZBLA 1Er3þ– 4Ce3þ PZG 0.5 Nd3þ ZBLA 1 Nd3þ

10 2 mm2 (charged) 10 5 mm2 (channel) 85 2 mm2 (charged) 40 ? mm2 (channel)

3.3

235 @ 980 nm 240 @ 980 nm >1000 @ 795 nm >1000 @ 798 nm

1.9 (7.0) 9.4 (10.7) 0.37 (0.45) 0.39 (0.48)

>0.5

[42]

<0.3

[42]

6

[43]

4.6

[43]

1.05

3.9 3.9 2.7

It is known that pump absorption is strongly enhanced at 980 nm with Yb3þ codoping. To take advantage of this in fluoride hosts, Er3þ/Yb3þ/Ce3þ tridoping is necessary to prevent upconversion of Er3þ by empting the 4I11/2 level. With this tridoping, the simulated gain could reach nearly 10 dB, even for waveguides as short as 4–5 cm [39] thus coming near the best performance reported on exchanged phosphate glass (i.e. 13 dB for a waveguide 5.5 cm long [40]). To emphasize the quality of these structures, it is worth recalling the performance of a ZBLAN fibre (33 dB for a fibre 10 m long [41]). 11.4.2

Lasers

Turning now to lasers, the only demonstration of laser action comes from a fluoride glass channel waveguide obtained by direct UV writing of a negative index change close to 0.01 in a fluoro-aluminate glass layer fabricated by dip casting [38]. The negative index change induced by photothermal expansion of the glass produces lateral confinement. Table 11.4 gives the characteristics of this laser and compares them to oxide glass waveguides. As can be seen, the performances are not as good but they could be largely improved considering the highly multimode structure. Green upconversion lasing was unsuccessful in a ridge ZBLAN waveguide obtained by pulse laser deposition and laser micromachining [30]; possible reasons are large scattering and coupling loss inside the resonator. However, simulation predicts that a conversion efficiency of 13 % is possible with a loss by scattering reduced down to 1 dB/cm (the actual value is 5 dB/cm).

342

Functionalized Inorganic Fluorides

It is clear that the results are not as good as expected, considering the intrinsic advantage of fluoride in comparison to oxide glasses. Unfortunately, the channel waveguide with the lowest propagation loss (i.e. 0.1 dB/cm) is achieved by lateral confinement for which control of the waveguide width (10 mm) is not as accurate as for the photolithography process.

Table 11.5 performance of fluoride glass waveguide lasers at 1.05 and 1.3 mm (pumping @ 800 nm). The oxide waveguides exhibiting the best performance are given for comparison l (mm)

Glass

WG section (type)

slope efficiency

1.05

ALF70 Nd3þ BK7* Nd3þ ALF70 Nd3þ SiO2-P2O5 Nd3þ

10 100 mm2 (channel) 4 5 mm2 (channel) 10 100 mm2 (channel) 8 12 mm2 (ridge)

27%

1.3

power threshold

length (cm)

loss (dB/ cm)

Ref

60 mW

1.6

0.1

[38]

42%

10 mW

2.4

0.2

[44]

2%

32 mW

1.6

0.1

[38]

2.6%

20 mW

5.9

<0.8

[45]

* BK7: exchanged Naþ/Kþ borosilicate glass

11.5

Fluoride Transparent Glass Ceramics: An Emerging Material

A new kind of material doped with rare-earth ions have been intensively investigated in the last decade – transparent glass ceramics (GC). These materials are obtained by controlled crystallization of a fraction of glasses by thermal process to get the active ions embedded in the crystal phase (see also Chapter 9 in the present book). Glass ceramics are thus of great importance in photonics because they can offer higher cross sections compared to glass that can be exploited in order to fabricate more compact devices. While the spectroscopic efficiency of GCs has been demonstrated in bulk, published works on rare-earth-doped GC waveguides remain very few, even for oxide materials [46]. The major obstacle concerns the crystallite size which has to be sufficiently small (a few nanometers) to keep Rayleigh scattering losses to an acceptable low level. Jestin et al. succeeded with SiO2-HfO2 glass made by sol-gel with propagation loss around 1 dB/cm after thermal treatment [47]. The first rare-earth -doped transparent CG containing fluoride was reported by Auzel et al. [48], combining pure oxide components with PbF2. A few years later, they obtained a totally fluoride transparent GC starting from ZELA or ZELAG glasses [49, 50] (see Table 11.1 for composition). The particularity of these glasses is to show only one crystalline phase through a spinodal decomposition, unlike ZBLAN glass, which gives nontransparent GC. It is also worth noticing that this glass is not significantly affected by concentration quenching up to 10 mol%. Recently, ZELA glass waveguides have been successfully obtained by PVD, through coevaporation of a LaF3-ErF3 mixture and ZBNA (ZrF4-BaF2-NaF-AlF3) glass [22]. Figure 11.5 illustrates the process by showing the vapour pressure curves of the fluoride compounds involved in the system. With a suitable thermal treatment, vitreous

Fluoride Glasses and Planar Optical Waveguides ZrF4

4

GaF3 AlF3 NaF

2 0 log(P Torr)

343

BaF2 ErF3 LaF3

–2 –4 –6 –8 –10 500

Figure 11.5

700

900 1100 Temperature (°C)

1300

Vapour pressure curves of fluorides entering ZELAG et ZBNA glass compositions

Normalized intensity

thin films deposited on a substrate heated slightly above Tg crystallize in a single phase with composition close to the glass one, in a similar way to the bulk (two-step process). When increasing deposition temperature, GC containing LaF3 nanocrystals (doped with 30 mol% Er3þ) are obtained without any thermal treatment (one-step process)

1.0

GC glass

0.8

crystal LaF3 : 0.3Er3+

0.6 0.4

LaZr3F15 : Er3+

0.2 0.0

one-step process

–0.2 –0.4 two-steps process –0.6 1400

1450

1500 1550 1600 Wavelength (nm)

1650

1700

Figure 11.6 Luminescence spectra of Er3þ at 1.5 mm upon 514.5 nm excitation of ZELA waveguides containing either LaF3 (one-step process) or LaZr3F15 (two-steps process) nanocrystals. The spectra of the pure crystalline phases and glass are given for comparison

344

Functionalized Inorganic Fluorides

it is likely that small ‘aggregates’ of ErF3 and LaF3 formed in the vapour state and quenched on the substrate act as nucleation agents. Guiding properties in infrared are not affected by the presence of nanocrystals in the waveguide with respect to the glassy waveguide for which propagation loss is 1.3 dB/cm, consistent with the dimension of the nanocrystals (<50 nm). The changes in the Er3þ environment in the precursor glass and GC are visible on the luminescence spectra at 1.5 mm (Figure 11.6) with Er3þ both in glassy and crystalline parts [51]. For LaF3 containing GC this results in a flatter and broader emission band (71 nm at half height width) close to the highest value reported in oxyfluoride bulk GC containing PbF2 crystals [52]. Concerning emission efficiency, some more work is necessary to optimize Er3þ doping to avoid concentration quenching in the crystalline phase.

11.6

Conclusion

In this chapter, we have shown that fluoride glass-confined waveguides with an acceptable propagation loss level (i.e. <0.5 dB/cm) can be obtained by suitable combination of composition and fabrication technique. This was a real challenge knowing the low thermal stability of fluoride glasses toward crystallization in comparison to oxide homologues. Actually, the performance of planar optical waveguides is less impressive than in fluoride glass optical fibre amplifiers or lasers; it partially reflects the time taken to establish reliable processes before achieving the required guiding structure. However, it is hoped that the performance as well as the integration may be further improved by the increasing number of laboratories working in this field, driven by the need for low-cost and reliable devices, particularly in the fibre telecommunication industry. In this regard, the recently emerged fluoride glass-ceramic waveguide appears as a promising structure in terms of bandwidth and flatness of the Er3þ emission band at 1.55 mm for optical amplification. Rare-earth doped glass waveguides may also find applications in realizing efficient compact visible laser sources, using upconversion processes by energy transfer, favoured by high doping rate and/or codoping.

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[31] E. Lebrasseur, Y. Gao, B. Boulard, B. Jacquier, Amplification in Er3þ doped PZG fluoride glass channel waveguides, ECOC Proceeding, Nice, paper I-54, (1999). [32] Y. Gao, O. Perrot, B. Boulard, J.E. Broquin, R. Rimet and C. Jacoboni, Preparation of PZG fluoride glass channel waveguides, J. Non-Cryst. Solids, 213&214, 137–140 (1997). [33] J.D. Shephard, D. Furniss, P.A. Houston, A.B. Seddon, Fabrication of mid-infrared planar waveguides from compatible fluorozirconate glass pairs, via hot spin-casting, J. Non-Cryst. Solids, 284, 160–167 (2001). [34] J. Lousteau, D. Furniss, A.B. Seddon, P. Sewell and T.M. Benson, Fluoride glass planar waveguides for active applications, Mat. Sci. and Engineering B, 105(1–3), 74–78 (2003). [35] Y. Kawamoto, M. Teramoto, T. Hatano, M. Shojiya, Preparation and etching processing of planar thin film of Pr3þ-doped fluorozirconate glass, J. Mater. Sci., 20, 5013–5016 (2001). [36] K. Miura, J. Giu, T. Mitsuyu, K. Hirao, Preparation and optical properties of fluoride glass waveguides induced by laser pulses, J. Non-Cryst. Solids, 256&257, 212–219 (1999). [37] B. Boulard, L. Brilland, H. Poignant, U.V. Writing of channel waveguides in erbium doped fluoride glass thin films, Electron. Lett., 34(3), 267–268 (1998). [38] D.W. Harwood, A. Fu, E.R. Taylor, R.C. Moore, Y.D. West, D.N. Payne, A 1317nm neodymium doped fluoride glass waveguide laser, ECOC Proceeding, Munich-Germany, 2, 191–192 (2000). [39] J.-L. Adam, Compositions, properties and applications of non-oxide glass, in in Photonic Glasses, Research Signpost, Editor R. Balda, 45–65(2006). [40] J.M.P Delavaux, G.C. McIntosh, G.R. Shmulovich, J. Kerkovian, A. Barbier, Multiple carrier analog transmission system with Er3þ doped planar optical waveguide amplifiers, Proc. OFC, 4, 64–66 (2000). [41] Y. Miyajima, T. Komukai, T. Sugawa and T. Yamamoto, Rare earth-doped fluoride fibre amplifiers and fibre lasers, Opt. Fibre Technol., 1, 35–47 (1994). [42] I. Vasilief in Guide d’onde canaux amplificateurs en verres de fluorures dope´s erbium: spectroscopie et amplification optique, PhD thesis, Universite´ Lyon (2003). [43] E. Lebrasseur, B. Jacquier, M.C. Marco de Lucas, E. Josse, J.-L. Adam, G. Fonteneau, Y. Gao, B. Boulard, C. Jacoboni, J.E. Broquin and R. Rimet, Optical amplification and laser spectroscopy of neodymium doped fluoride glass channel waveguides, J. Alloys and Compounds, 275–277, 716–720 (1998). [44] D.P. Shepherd, S.J. Hettrick, C. Li, J.I. Mackenzie, R.J. Beach, S.C. Mitchell, H.E. Meissner, High-power planar dielectric waveguide lasers, J. Phys. D: Appl. Phys., 34, 2420–2432 (2001). [45] J.R. Bonar, J.A. Bebbington, J.S. Aitchison, G.D. Maxwell, B.J. Ainsly, Low threshold Nddoped silica planar waveguide laser, Elect. Lett., 30(3), 229–230 (1994). [46] V.K. Tikhomirov, A.B. Seddon, J. Koch, D. Wandt and B.N. Chichkov, Fabrication of buried waveguides and nanocrystals in Er3þ-doped oxyfluoride glass, Phys. Stat. Sol.(a), 202(7), R73–R75 (2005). [47] Y. Jestin, C. Armellini, A. Chiasera, A. Chiappini, M. Ferrari, E. Moser, R. Retoux, Low-loss optical Er3þ-activated glass-ceramics planar waveguides fabricated by bottom-up approach Appl. Phys. Lett., 91, 071909, 1–3 (2007). [48] F. Auzel, D. Pecile and D. Morin, Rare Earth Doped vitroceramics: new, efficient, blue and green emitting materials for infrared up-conversion, J. Electrochem. Soc., 122, 101–107 (1975). [49] F. Auzel, K.E. Lipinska-Kalita, P. Santa-Cruz, A new Er3þ-doped vitreous fluoride amplification medium with crystal-like cross-sections and reduced inhomogeneous line width, Opt. Mater., 5, 75–78 (1996). [50] M. Mortier, A. Monte´ville, G. Patriarche, G. Maze´, F. Auzel, New progresses in transparent rare-earth doped glass-ceramics, Opt. Mater., 16, 255–267 (2001). [51] B. Boulard, O. Pe´ron, Y. Jestin, M. Ferrari, C. Duverger-Arfuso, Y. Gao, Characterizations of Er3þ-doped fluoride glass ceramics waveguides containing LaF3 nanocrystals, to be published in J. Luminescence, 129, 1637–1640 (2009). [52] V.K. Tikhomirov, D. Furniss, A. Seddon, I.M. Reaney, M. Beggiora, M. Ferrari, M. Montagna and R. Rolli, Fabrication and characterization of nanoscale, Er3þ-doped, ultratransparent oxyfluoride glass ceramics, App. Phys. Lett., 81, 1937–1939 (2002).

12 Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates Karim Adil, Amandine Cadiau, Annie He´mon-Ribaud, Marc Leblanc and Vincent Maisonneuve Laboratoire des Oxydes et Fluorures, UMR 6010 CNRS, Faculte´ des Sciences et Techniques, Universite´ du Maine, Avenue O. Messiaen, 72085 Le Mans Cedex 9, France

12.1

Introduction

Depending on the acid-base properties of metal cations in solutions, the condensation of cationic or anionic metal species follows different paths to give crystallized solids. In water solutions, hydration, olation, oxolation and complexation mechanisms are involved in the formation of initial nuclei and the growth or the ageing of solid particles[1, 2]. Such mechanisms depend strongly on the concentrations of the reactants and the resulting pH of the solutions, on the temperature, and, to a lesser degree, on the nature of counter cations. They are well documented for oxides or ‘ate’ salts (silicates, phosphates, vanadates, tungstates. . .)[3]; numerous compounds are described[4, 5] and lead to industrial materials. On the other hand, the literature on fluoride crystallization is scarce, in spite of the recent work on nanosized fluoride particles for applications in optics[6] or catalysis[7]. Interest in fluorides is mainly driven by their properties as crystallization agents for the synthesis of zeolites[8] or phosphate-based materials[9] and more recently of metal organic frameworks[10] (MOF). Mixed anions (fluoride-phosphate, fluoride-carboxylate . . .) build three-dimensional 3D open frameworks for gas storage (CO2), catalysis or emerging applications[11,12] (the symbols 0D, 1D, 2D, 3D, characteristic of the dimensionality of the frameworks, are used throughout this chapter).

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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Functionalized Inorganic Fluorides

Fluoride anion architecture is only encountered in metal fluorides templated by amine cations (the amine name abbreviations are listed in Table 12.1). Zirconates[13–15] are extensively studied together with aluminates[16, 17], beryllates[18], scandates[19], tantalates [20, 21], titanates[22, 23] and uranates[24]. Very few compounds present 3D frameworks: only one beryllate, [Hmeam] (Li2Be2F7)[25], two yttriates, [H2en]0.5 (Y2F7)[26] and [H2dap]0.5 (Y2F7)[27], and one zirconate, (H3O) [Hgua]5 (ZrF5)6[28]. The largest pores are found in this last fluorozirconate. Table 12.1 col dado dap dien en gua guan guaz meam phenetam pipz py tren triaz 4-amtriaz 5-amtetraz

Amine name abbreviations 2,4,6-trimethylpyridine ¼ collidine 1,12-diaminododecane 1,4-diaminopropane bis-(2-aminoethyl)amine 1,2-diaminoethane guanidine guanine 3,5-diamino-1,2,4-triazole ¼ guanazole methylamine 1-phenylethylamine piperazine pyridine tris-(2-aminoethyl)amine 1,2,4-triazole 3-amino-1,2,4-triazole 5-amino-tetrazole

CH3(C5H2N)(CH3)2 H2N(CH2)12NH2 H2N(CH2)3NH2 (NH2(CH2)2)2NH H2N(CH2)2NH2 HN ¼ C(NH2)2 H2N(C2HN3)NH2 NH2CH3 (C6H5)(CH2)2NH2 (H2N(CH2)2)3N

C8H11N C12H28N2 C3H10N2 C4H13N3 C2H8N2 CH5N3 C5H5N5O C2H5N5 CH5N C8H11N C4H10N2 C5H5N C6H18N4 C2H3N3 C2H4N4 CH3N5

This chapter is devoted to the crystallization of fluoroaluminates. Emphasis is given to hybrid organic-inorganic salts that are synthesized in hydro(solvo)thermal conditions; numerous polyanions, derived from the primary building anion (AlF6)3, are now evidenced, together with extended one- or two-dimensional architectures. These species are compared with the (Al, F) units that are found in inorganic fluoroaluminates for which a classification of minerals was proposed by F. C. Hawthorne in 1984[29]. The condensation of the AlF6 units to give polyanions results from the elimination of one or two fluorine atoms and can be described from the association of the AlF6 octahedra by one corner or by one edge (the association of two octahedra by one face has not yet been found in hybrid fluorides). As a consequence, the average Al-F distances fall into three ˚ for nonbridging fluorine atoms, 1.82 A ˚ for corner bridging atoms and 1.88 A ˚ types: 1.80 A for edge-bridging atoms; moreover, the distortion of the AlF6 octahedra is the smallest for isolated octahedra and the largest for edge-connected octahedra. The stability of the hybrid organic-inorganic structures is partly due to the hydrogen bonds that link the hydrogen atoms of amine groups to fluorine atoms.

12.2

Synthesis

Inorganic fluorides can be prepared in the solid state at high temperature (ceramic route). Anhydrous conditions must prevail because of the sensitivity of fluorides towards

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

349

hydrolysis. The starting fluorides are generally heated in sealed gold or platinum tubes for 1–3 days to obtain polycrystalline powders[30]. Inorganic crystals grow in chloride fluxes that behave as good solvents for numerous fluorides[31]. After dissolution of the starting fluorides in the chloride melt, the temperature is slowly decreased (0.5–5 C/min) down to the eutectic temperature. Most often, chlorides in the liquid state prevent fluorides from hydrolysis. In 1975 it was shown, unexpectedly, that anhydrous ferric fluoride crystals grow in concentrated HF supercritical solutions[32] (380 C, 200 MPa, 10 days) and numerous phases were further obtained: low temperature phases[33], metastable phases[34], mixed valence phases or hydrates[35]. . . Later, subcritical conditions (180–220 C, 0.1–0.3 MPa) were found to promote the crystallization of oxide fluorides[36]. Recently, microwave heating of HF solutions in water or ethanol was proved to be efficient for the elaboration of hybrid organic-inorganic fluorides[37]. Single crystals, suitable for X-ray diffraction data collection, grow in spite of the very short duration of the syntheses (30–120 min). The exploration of large concentration domains of the starting soluble materials is then possible[38]. It is convenient to give composition space diagrams in order to establish the crystallization domains of the precipitates[39, 40]. These diagrams are represented at a given concentration of metal cation source (aluminum hydroxide in Figure 12.1) and a given temperature; the axes give the ratio of the starting materials. Figure 12.1 gives the example of the Al(OH)3-guaz-HFaq.-EtOH system[41]. Four zones correspond to the crystallization of three organic-inorganic hybrid fluoroaluminates and one inorganic fluoride (Al2[F0.5(OH)0.5]6 H2O pyrochlore). The position of the dots gives the composition of the starting mixture and the associated symbol indicates the nature of the solid precipitate.

guaz 50 50 [[Hguaz]2·(AlF5(H2O))·2H2O 60

[Hguaz]2·(AlF5)

70 80

40 30

[Hguaz]2·(Al2F8)

20

90 [Hguaz]2·(Al5F17)

10 Al2(F0.5,OH0.5)6·H2O

100 HF 0

10

20

30

40

0 50 Al(OH)3

Figure 12.1 Composition space representation of the Al(OH)3-guaz-HFaq.-EtOH system for [Al3þ] ¼ 1 mol.L1 and T ¼ 190 C

350

12.3

Functionalized Inorganic Fluorides

Extended Finite Polyanions (0D)

Aluminum cations Al3þ adopt preferentially the octahedral coordination in hybrid fluorides and a tetrahedral coordination is only encountered in four compounds, [N(CH3)4] (AlF4)[42, 43], [HN2C14H18] (AlF4)[44] and [Hcol]6 (AlF4)3 (Al3F12)[44] (associated with Al3F12 chains, see 12.4.5.5) and, [RePO3NC20H20] (AlF4)[45]. Recently, a fivefold coordination was suggested by NMR spectroscopy in [N(CH3)4]2 (AlF5)[46]; however, the structure is still unknown. A 0D network of AlF6 octahedra is found for F/Al ratios in the range 6 £ F/Al £ 35/8. 12.3.1

Isolated AlF4 Tetrahedra

Fifteen years ago, N. Herron et al.[44] succeeded in preparing, in anhydrous conditions, the first aluminum fluoride with a fourfold coordination of Al3þ: [HN2C14H18] (AlF4) (N2C14H18 ¼ 1,8-bis(dimethylamino)naphthalene). The synthesis was performed in two steps: a pyridinium aluminate, [Hpy] (Al3F10), was first recovered from the reaction of trimethylaluminum and the (HF)x pyridine adduct in pyridine; then, the exchange of pyridinium cation with [HN2C14H18]þ was achieved in acetonitrile to give [HN2C14H18] (AlF4) (Figure 12.2b). Single crystals were further recrystallized in methylene chloride/toluene. The crystal structure determination indicated that one triangular face of the AlF4 tetrahedra is disordered on two pseudo-mirror symmetry-related positions (Figure 12.2a). a

b

c b

Figure 12.2 (a) Disordered AlF4 tetrahedron and (b) [100] projection of the structure of [HN2C14H18] (AlF4)

12.3.2

Isolated AlF6 Octahedra

Numerous AnA’mAlF6 compounds (A ¼ alkaline, A’ ¼ alkaline-earth) are reported in the literature and most of the natural fluoride solids found in the earth mantle are fluoroaluminates. Several minerals are well known archetypes of fluoride structures: Na3AlF6[47] cryolite, K2NaAlF6[48] elpasolite, Na2MgAlF7 [49] weberite, LiCaAlF6 [50] colquiriite, Na3Li3Al2F12[51] cryolithionite, NaCaAlF6 H2O thomsenolite (C 2/c)[52], and pachnolite (P2/c)[53]. Other simple archetypes are found in synthetic fluoroaluminates: NaSrAlF6[54], NaCaAlF6 ( and forms)[55, 56], Na2Ca3Al2F14[57].

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

351

Two common minerals only, cryolite and elpasolite, will be described, together with synthetic NaSrAlF6 and hybrid (H3O) [H2pipz] (AlF6). 12.3.2.1

Na3 AlF6 Cryolite

The cryolite structure derives from the structure of the orthorhombic perovskite GdFeO3[58]. The structure was determined by Hawthorne et al.[59] in 1975 (monoclinic, P21/n). The aluminum atoms of the isolated AlF6 octahedra are the nodes of a bodycentred I network (Figure 12.3); sodium atoms adopt a fourfold and sixfold coordination.

0

50

b

Na a

Figure 12.3 [001] Projection of the structure of Na3AlF6 cryolite

12.3.2.2

K2NaAlF6 Elpasolite

K2NaAlF6 elpasolite, cubic[48] is a perovskite superstructure with aelpasolite ¼ 2aperovskite. The AlF6 octahedra are centred at the origin and the face centres of the cubic cell (Figure 12.4). Sodium atoms are localized at the middle of the cell edges while potassium atoms lie at equal distances from four AlF6 octahedra.

50

25 b

0

0

0

Na

K 75 50

a

Figure 12.4 [001] Projection of the structure of K2NaAlF6 elpasolite

352

Functionalized Inorganic Fluorides

12.3.2.3

NaSrAlF6

Three phases, NaSrAlF6 [54], Na2Sr7Al6F34[60] and Na3Sr4Al5F26[61], are found in the ternary system NaF-SrF2-AlF3. NaSrAlF6, discovered in 1990, is the archetype of a new structural family that includes NaBaAlF6[62]. The isolated AlF6 octahedra lie in (100) planes at z ¼ 0.24, 0.4, 0.74 and 0.90; strontium and sodium atoms adopt an eightfold coordination (Figure 12.5). a c 24

25 74

76

Na 74

75 24

Sr 26

Figure 12.5

12.3.2.4

[010] Projection of the structure of NaSrAlF6

(H3O) [H2pipz] (AlF6)

The first hybrid fluoroaluminate, [Hgua]3 (AlF6), was evidenced by Bukovec[63] in 1983. This pioneering work inspired several other groups[16, 64–68] and several fluorides with isolated AlF6 octahedra were obtained. (H3O) [H2pipz] (AlF6) synthesized by Bentrup in 1996[65], can be described. The monoclinic structure (P21/c) is built up from two types of AlF6 octahedra, surrounded by three piperazinium cations and three hydronium H3Oþ ions (Figure 12.6).

a c

Figure 12.6 [010] Projection of the structure of (H3O) [H2 pipz] (AlF6)

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

12.3.3

353

Al2F11 Dimers

K3Ba7Al6F33Cl2[69] is the unique fluorochloroaluminate in which Al2F11 dimers are reported (Figure 12.7b). These Al2F11 dimers result from the connection of two AlF6 octahedra by one corner (Figure 12.7a). A disorder of fluorine atoms implies two possible orientations of the Al2F11 dimers. a

b

Cl Ba K b a

Figure 12.7

12.3.4

(a) Al2F11 dimer and (b) [001] projection of the structure of K3Ba7Al6F33Cl2

Al3F16 Trimers

Three AlF6 octahedra are associated by corners in Al3F16 trimers (Figure 12.8a). This Al3F16 unit is only found in Na2Sr7Al6F34[60] (Figure 12.8b) isostructural with Na2Sr6MgAl6F32(OH)2 jarlite[70–73]. a

b

b a

F Na

Sr

Figure 12.8 (a) Al3F16 trimer and (b) [001] projection of the structure of Na2Sr7Al6F34

12.3.5

Al2F10 Dimers

Two edge-sharing AlF6 octahedra build the Al2F10 dimers found in [Hpy]4 (Al2F10) 4H2O[74] (Figure 12.9a) and [H2pipz]2 (Al2F10) 2H2O[75]. In [Hpy]4 (Al2F10) 4H2O, the (Al2F10)4 anions lie in (100) planes with water molecules

354

Functionalized Inorganic Fluorides

(Figure 12.9b). These layers are separated by piperazinium cations. In [H2pipz]2 (Al2F10) 2H2O, the inorganic layers are similar; their stacking occurs along [101]. a

b

c b

Figure 12.9 (a) Al2F10 dimer and (b) [100] projection of a part of the structure of [Hpy]4 (Al2F10) 4H2O at 0 < x < ¼

12.3.6

Al4F20 Tetramers

Domesle et al. obtained Ba3Al2F12 in 1982. The orthorhombic structure is made up of Al4F20 [76] tetramers (Figure 12.10a). In Ba3Al2F12, these planar cycles of four corner sharing octahedra are centred at the nodes of a centred I network (Figure 12.10b). a

a

b c

Figure 12.10

12.3.7

Ba F

(a) Al4F20 tetramer and (b) [010] Projection of the structure of Ba3Al2F12

Al4F18 Tetramers

The highly symmetrical association of four corner sharing AlF6 octahedra builds tetrahedral Al4F18 units (Figure 12.11a) that are found in three hybrid fluoroaluminates: [H3dien]2 (Al4F18) [39] (Figure 12.11b), (H3O)2 [Hgua]16 (Al4F18)3 H2O [77] and [H3tren]2 (Al4F18) 3.5H2O [78].

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates a

355

b a c

c

d c

b

Figure 12.11 (a), (c) Al4F18 tetramers and projections of the structures of (b) [H3dien]2 (Al4F18) and (d) [H3tren]2 (Al4F18)

Both corner and edge-sharing associations of AlF6 octahedra are found in Al4F18 units (Figure 12.11c) of [H3tren]2 (Al4F18)[78] (Figure 12.11d).

12.3.8

Al5F26 Pentamers

The planar association by corners of four AlF6 octahedra to one central AlF6 octahedron builds Al5F26 pentamers (Figure 12.12a). In Na3Sr4Al5F26[61], these Al5F26 pentamers, related by a 42 helical axis, lie in perpendicular (a,c) and (b,c) planes at the origin and the face centres of the tetragonal cell (Figure 12.12b).

12.3.9

Al7F30 Heptamers

In 2002, Goreshnik et al. prepared (H3O) [H4tren]2 (Al7F30)[79] (Figure 12.13b). This fluoroaluminate of tren is built up from isolated (Al7F30)9 polyanions (Figure 12.13a). In (Al7F30)9, the association of seven corner-sharing octahedra can be described as two tetrahedra of octahedra sharing a common octahedron. In (H3O) [H4tren]2 (Al7F30), tren amines are tetraprotonated and adopt a ‘scorpion’-type configuration; H3Oþ cations ensure the charge compensation.

356

Functionalized Inorganic Fluorides a

c

b b

Sr Na

Figure 12.12

(a) Al5F26 pentamer and (b) projection of the structure of Na3Sr4Al5F26

a

b

a c

Figure 12.13 (a) Al7F30 (H3O) [H4tren]2 (Al7F30)

12.3.10

heptamer

and

(b)

projection

of

the

structure

of

Al8F35 Octamers

In [H3tren]4 (Al8F35) (OH) H2O[78], two tetrahedral Al4F18 units, linked by one corner, form the largest known fluoride polyanion, (Al8F35)11 (Figure 12.14a). In this phase, the presence of OH hydroxyl groups, surrounded by amine cations, is not clearly ascertained (Figure 12.14b). 12.3.11 12.3.11.1

Mixed Polyanions AlF6 and Al2F11 Polyanions

Both AlF6 and Al2F11 units are found in [H3tren]4 (AlF6)2 (Al2F11) (F) 10H2O[78]. One ‘free’ fluoride anion is octahedrally coordinated by six hydrogen atoms of NH3 groups from two amine cations (Figure 12.15a). H2O molecules, at z ¼ and z 3⁄4, build hexagonal rings (Figure 12.15b) that are inserted between the neutral [H3tren]4 (AlF6)2 (Al2F11) (F) layers (Figure 12.15c).

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates a

357

b OH c

Ow

b

OH Ow

Figure 12.14 (a) Al8F35 octamer [H3tren]4 (Al8F35) (OH) H2O a

and

(b)

projection

of

the

structure

of

b

F

b a c

c b

F Ow

F

Figure 12.15 (a) ‘‘Free’’ fluoride anion environment, (b) water layer at z ¼ and (c) [100] projection of the structure of [H3tren]4 (AlF6)2 (Al2F11) (F) 10H2O

358

Functionalized Inorganic Fluorides

12.3.11.2

Al2F11 and Al4F18 Polyanions

In [Hphenetam]11 (Al2F11) (Al4F18) 7H2O[80], corner sharing Al2F11 dimers and tetrahedral Al4F18 tetramers lie, with water molecules in (001) layers (Figure 12.16). The amines, on both sides of the (Al2F11) (Al4F18) 7H2O layers, orientate the –NH3 groups towards the preceding layers and the planes of phenyl cycles are approximately perpendicular to (001). Consequently, the [Hphenetam]11 (Al2F11) (Al4F18) 7H2O sheets are only linked by weak Van der Waals bonds.

b a

Figure 12.16

12.4

Al2F11 dimers and Al4F18 tetramers in [Hphenetam]11 (Al2F11) (Al4F18) 7H2O

1D Networks

Simple chains or ramified chains are found in fluoroaluminates and the AlF6 octahedra are linked either by corners or by edges. The formulation of the inorganic network ranges from AlF5 to AlF4.

12.4.1 12.4.1.1

AlF5 Chains Trans Connection

Corner sharing of AlF6 octahedra by opposite vertices built infinite trans chains AlF5 (Figure 12.17a). In -CaAlF5[81, 82], the AlF5 chains lie along the [001] direction (Figure 12.17b).

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates a

359

b

b Ca

a

Figure 12.17 (a) trans-chain AlF5 and (b) [001] projection of the structure of -CaAlF5

12.4.1.2

Cis Connection

In cis chains AlF5, the AlF6 octahedra share adjacent corners (Figure 12.18a). Such chains were first found in -BaAlF5[83] and, later, in other , g, d varieties of BaAlF5 [82, 84]. These forms differ by the distortions of the chains. Recently, a series of hybrid fluorides with cis-AlF5 chains was evidenced: [H3N(CH2)xNH3] (AlF5)[85] (x ¼ 6, 8, 10, 12) (Figure 12.18b). In this series, the AlF5 chains are very similar. a

b

a c

Figure 12.18

12.4.2

(a) cis-chain AlF5 and (b) [010] projection of the structure of [H2dado] (AlF5)

Al2F9 Chains

The decomposition of K(H3O)AlF6[86–88] crystals at room-temperature is slow and gives K2(H5O2)Al2F9[89]. In this last phase, the double Al2F9 chains are described from the association by corners of two simple trans-AlF5 chains (Figure 12.19a). The double chains are separated by potassium and (H5O2)þ cations (Figure 12.19b). a

b

O

O K

b a

Figure 12.19

(a) Al2F9 chain and (b) [001] projection of the structure of K2(H5O2)Al2F9

360

Functionalized Inorganic Fluorides

12.4.3

Al7F29 Chains

Al7F30 heptamers can be associated by opposite corners in order to build infinite chains Al7F29[79] in [H4tren]2 (Al7F29) 2H2O (Figure 12.20a). Tren cations adopt a ‘scorpion’ type configuration (Figure 12.20b). a

b a c

Figure 12.20 (a) Al7F29 chain and (b) [010] projection of the structure of [H4tren]2 (Al7F29) 2H2O

12.4.4

AlF4 Chains

Three association modes of AlF6 octahedra to give double, triple or ramified AlF4 chains are reported. Double chains involve both corner and edge sharing while corner sharing is only found in triple and ramified chains. Al2F8 Chains

12.4.4.1

Until 2004, hybrid fluoroaluminates were always prepared with organic (poly)amines. Recently, Loiseau et al.[90] evidenced [Hpy]2 [C6H3(CO2H)3] (Al2F8) where pyridinium cations and carboxylic acid, benzene 1,3,5 tricarboxylic acid, are associated (Figure 12.21b). The double chains Al2F8, also found in [Hguaz] (Al2F8)[41] and in a new variety of KAlF4 [91] , result from the connection by edges of trans-AlF5 chains (Figure 12.21a). a

b

c b

Figure 12.21 (a) Al2F8 chain [Hpy]2 [C6H3(CO2H)3] (Al2F8)

and

(b)

[100]

projection

of

the

structure

of

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

12.4.4.2

361

Al3F12 Chains

In 1979, simultaneously, Fourquet et al.[92] and Lo¨sch et al.[93] published the synthesis and the structure of -CsAlF4. Later, single crystals were grown in HF solution using a hydrothermal method in supercritical conditions[94] (P ¼ 2500 bar, T ¼ 500 C). In -CsAlF4, trans-AlF5 chains are associated by adjacent corners to build infinite triple chains Al3F12 (Figure 12.22a). These [0001] chains are centred at the cell corners and separated by Csþ cations (Figure ˚ . A similar arrangement of the Al3F12 chains is found 12.22b); the interchain distance is 9.49 A ˚. in [Hguan]3 (Al3F12)[95]; the interchain distance is enlarged to 14.2 A a

b

b

Cs a

Figure 12.22 (a) Al3F12 chain and (b) [001] projection of the structure of -CsAlF4

12.4.4.3

Al6F24 Chains

In 1992, Fourquet et al. studied the transformation of caesium fluoroaluminates, eventually hydrated, with the temperature and with the nature of the atmosphere. Under neutral atmosphere, the following reaction, leading to a new form of CsAlF4, was 150C observed[92, 93, 96]: Cs[AlF4(H2O)2] ! g-CsAlF4. In g-CsAlF4, the monoclinic structure (Figure 12.23a) contains infinite grafted chains that can be considered as trans-connected AlF5 chains decorated by Al2F11 dimers. In this chain, the aluminum atoms lie at the corners of vertex sharing Al4 tetrahedra (Figure 12.23b). This chain is also described from (Al6F27)9 or (Al7F29)8 anions; the Al6F27 units are linked by corners to give [Al6F21F6/2] or Al6F24 chains, while the Al7F29 entities share two opposite AlF6 octahedra. Recently, similar chains were found in the hybrid fluoroaluminates (H3O) [H2pipz] (Al3F12)[97] (Figure 12.23c) and [H4tren]3/2 (Al6F24) 3H2O[98]. They were obtained, respectively, from AlF3 and piperazine or Al(OH)3 and tren in HF solutions and hydrothermal conditions. 12.4.5

Mixed Polyanions and/or Chains

All chains are built up from the association of AlF6 octahedra whereas polyanions are either AlF6 octahedra or AlF4 tetrahedra. 12.4.5.1

AlF6 Octahedra and trans-AlF5 Chains [99]

Pb5Al3F19 results from the reaction of PbF2 and AlF3 in the solid state. Four phase transitions occur at 160, 295, 320 and 370 K[100]. All varieties contain trans-AlF5 chains and isolated AlF6 octahedra (Figures 12.24a and b). They differ by the relative rotations or

362

Functionalized Inorganic Fluorides b

a c

c

a

b

c

Figure 12.23 (a) Projection of the structure of g-CsAlF4 along [010], (b) Al3F12 chain and (c) projection of the structure of (H3O) [H2pipz] (Al3F12) along [100]

tilting of the AlF6 octahedra and AlF5 chains. The coordination of the Pb2þ ions, from six to ten, includes fluorine atoms from the AlF6 units or AlF5 chains and extra ‘free’ fluorine atoms (Figure 12.24c).

a

b

c

b c

a

F

Pb

Figure 12.24 (a) trans-chain AlF5, (b) isolated AlF6 octahedron and (c) projection of the structure of Pb5Al3F19

12.4.5.2

AlF6 Octahedra and cis-AlF5 Chains

Ba3Cu2Al2F16, evidenced by Gredin et al. in 2003[101], presents infinite cis-AlF5 chains and isolated AlF6 octahedra (Figures 12.25a and b). Copper atoms, in octahedral coordination, connect the AlF6 octahedra while barium atoms link the cis-AlF5 chains to the copper aluminum fluoride layer (Figure 12.25c). 12.4.5.3

Al2F10 Dimers and Grafted Al3F15 Chains

˚ c ¼ 14.322(2) A ˚ ), isolated Al2F10 In one tetragonal form of I-SrAlF5[102] (a ¼ 19.882(2) A dimers (Figure 12.26b) and grafted Al3F15 chains (Figure 12.26a) are associated. The

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates a

363

c

F b

Cu

Ba c b

Figure 12.25 (a) cis-chain AlF5, (b) isolated AlF6 octahedron and (c) [100] projection of the structure of Ba3Cu2Al2F16

Al3F15 chains lie along a 42 axis that implies successive 90 rotations of the satellite octahedra (Figure 12.26c)

a

c

b

Sr b a

Figure 12.26 I-SrAlF5

12.4.5.4

(a) Al3F15 chain, (b) Al2F10 dimer and (c) [001] projection of the structure of

AlF5 Chains and Grafted Al3F15 Chains

˚ , c ¼ 7.617(2) A ˚ ) similar In a second tetragonal variety of II-SrAlF5[103] (a ¼ 14.089(2) A Al3F15 chains are found to be associated with trans-AlF5 chains (Figures 12.27a and b). The [001] projection of strontium atom positions are very similar in both tetragonal forms of SrAlF5 (Figure 12.27c).

364

Functionalized Inorganic Fluorides a

c

b

Sr

b a

Figure 12.27 II-SrAlF5

12.4.5.5

(a) trans-chain AlF5, (b) Al3F15 chain and (c) [001] projection of the structure of

AlF4 Tetrahedra and Al3F12 Chains

Tetrahedral and octahedral coordination of Al3þ is only reported in one collidinium compound, [Hcol]6 (AlF4)3 (Al3F12)[44] (Figure 12.28c). The AlF4 units (Figure 12.28a) are isolated while the AlF6 octahedra are associated by edges and corners to give infinite chains Al3F12 (Figure 12.28b). This phase is prepared from pyridinium aluminate, [Hpy] (Al3F10) (see 12.3.1 and 12.5.5), by a metathesis exchange reaction of pyridinium cation with [Hcol]þ in acetonitrile.

a

c

b

b a

Figure 12.28 (a) AlF4 tetrahedron, (b) Al3F12 chain and (c) [001] projection of [Hcol]6 (AlF4)3 (Al3F12)

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

12.5

365

2D Networks

The association of AlF6 octahedra to give infinite layered structures occurs for F/Al ratios in the range 14/3 £ F/Al £ 10/3.

12.5.1

Al3F14 Layers

The tetragonal layered structure of Na5Al3F14 chiolite[104] exhibits square cycles of eight corner-sharing octahedra (Figure 12.29a). Sodium cations are inserted in and between the (001) Al3F14 sheets, related by I centring (Figure 12.29b).

a

b

a c

Na

Na

b a

Figure 12.29

12.5.2

(a) Al3F14 sheet and (b) [010] projection of the structure of Na5Al3F14

AlF4 Layers

Like the AlF4 chains, three association modes of AlF6 octahedra build AlF4 layers. All modes are represented by inorganic fluorides. 12.5.2.1

Perovskite Layers

Simple perovskite-type layers were first evidenced by Brosset in KAlF4[105] in 1938 (Figure 12.30a); the symmetry is tetragonal and the Al-F-Al angle is ideal, 180. Numerous other layered fluorides were found further; however, most of the structures are distorted, due to a concerted rotation and/or a tilting of the octahedra. For example, the rotation of the AlF6 octahedra around the fourfold axis is 3.0 in – RbAlF4 [106] (Figure 12.30b). 12.5.2.2

TTB Layers

-RbAlF4, isostructural with -CsAlF4[96], presents TTB (tetragonal tungsten bronze)type AlF4 layers with a tetragonal symmetry. The corner-sharing octahedra leave triangular, square and pentagonal windows (Figure 12.31).

366

Functionalized Inorganic Fluorides a

b

b K

a

b a

Figure 12.30

Rb

[001] Projections of the structures of (a) KAlF4 and (b) -RbAlF4

Rb

b a

Figure 12.31 [001] Projection of the structure of -RbAlF4

12.5.2.3

HTB Layers

In Cs2NaAl3F12[107], the symmetry is rhombohedral (Figure 12.32b) and the HTB (hexagonal tungsten bronze)-type AlF4 layers exhibit triangular or hexagonal windows only. The AlF4 layer is described from the connection of Al3F15 trimers (Figure 12.32a). 12.5.3

Al2F7 Layers

Two preceding AlF4 layers of Cs2NaAl3F12 are linked by the axial corners of the octahedra in Rb2NaAl6F21[108] (Figure 12.33a). The resulting Al2F7 sheets stack along the c axis of the monoclinic cell (Figure 12.33b). Cs2NaAl3F12 and Rb2NaAl6F21 are part of the series A2Na(AlF3xþ1)3; the limit of this series is the 3D HTB type -AlF3 for x ! 1. 12.5.4

Al5F17 Layers

A new layer formulation was recently found in several hybrid fluoroaluminates: [Hgua]2 (Al5F17)[77] (Figure 12.34b), [Hguaz]2 (Al5F17), [H4-amtriaz]2 (Al5F17),

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

367

b

a a

b c

Cs Na

Cs

b c a Na

Figure 12.32

(a) HTB AlF4 sheet and (b) [112] projection of the structure of Cs2NaAl3F12

a

b

a c Na

Rb

b a

Figure 12.33 along [010]

Projections (a) of a double layer Al2F7 along [001] and (b) of Rb2NaAl6F21

[Htriaz]2 (Al5F17) and [H5-amtetraz]2 (Al5F17)[109]. The Al5F17 sheets can be described from the connection of Al4F16 columns of the ReO3 type structure with AlF5 chains or from the connection of HTB type Al5F19 columns or from the intergrowth of HTB type Al5F19 columns and ReO3 type Al4F16 columns (Figure 12.34a).

a

c

a

b

b

b a

Figure 12.34

(a) View of a layer Al5F17 and (b) projection along [001] of [Hgua] (Al5F17)

368

12.5.5

Functionalized Inorganic Fluorides

Al3F10 Layers

The thermal decomposition of [Hpy] (AlF4) at 275 C gives [Hpy] (Al3F10). The structure of [Hpy] (Al3F10) (Figure 12.35b) was only determined in 1999 after a ‘Dupont powder challenge’[110] is launched on a compound previously published by Herron et al.[111] as HAlF4. Chemical reanalysis, synchrotron and neutron powder data collection enabled the ab initio structure solution to be found. The structure consists of (001) Al3F10 layers (Figure 12.35a) separated by pyridinium cations. These layers can be considered as the association of Al3F14 trimers along b in order to build [010] Al3F12 chains that are further connected by corners along a.

a

b

a c

b a

Figure 12.35

12.6

(a) Al3F10 layer and (b) [010] projection of the structure of [Hpy] (Al3F10)

3D Networks

A 3D network of connected AlF6 octahedra with large open pores is still awaited. It can be expected that the stability of such an architecture is weak. Until now the species that can be inserted in small cavities are inorganic. 12.6.1

Al7F33 Network

Two fluoroaluminates of the ternary system NaF-CaF2-AlF3 are only reported: NaCaAlF6 and NaCaAl1.75F8.25[112] (or Na4Ca4Al7F33). The cubic structure of NaCaAl1.75F8.25, determined by He´mon et al. in 1990[56], is described from two interpenetrating cubic Al7F33 subnetworks (Figure 12.36b). Every corner of both cubes is decorated with one heptamer Al7F36 (one central octahedron and six corner sharing octahedra along –a, –b, –c) (Figure 12.36a). The heptamers are connected by corners along a, b, c. 12.6.2

Al2F9 Network

A 3D fluoroaluminate subnetwork, considered as lacunar HTB -AlF3, is found in KCaAl2F9[113]. One third of the aluminum cations of HTB layers are vacant in order to

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates a

369

b

c

a b

Figure 12.36

(a) Al7F36 heptamer and (b) cubic Al7F33 networks in Na4Ca4Al7F33

leave infinite trans-AlF5 chains. These chains, connected by corners along [001] (Figure 12.37b), run along [110] at z ¼ 1/12 and 7/12 (Figure 12.37a), along [110] at z ¼ 5/12 and z ¼ 11/12, along [010] at z ¼ 3/12 and 9/12. b

a

K Ca c a

Figure 12.37

12.6.3 12.6.3.1

b a

b

Projections of the structure of KCaAl2F9 (a) along [110] and (b) along c

AlF3 Network g-AlF3 (or Z-AlF3)

Thermal decomposition of [Hpy] (AlF4) gives the cubic variety g-AlF3[111], isostructural with pyrochlore: Al2[F0.5(OH)0.5]6 H2O [114]. The structure can be described from HTB layers connected by extra octahedra (Figure 12.38a and b) and from Al4F18 units linked by corners. Large tunnels with a hexagonal section run along [110] or [111] and in equivalent directions. 12.6.3.2

-AlF3

Thermal decomposition of -NH4AlF4[115] and dehydration of -AlF3 3H2O[116, 117] give the HTB -AlF3. The connected HTB AlF4 layers leave hexagonal tunnels along the c axis

370

Functionalized Inorganic Fluorides a

b

HTB

HTB

b a

c HTB

Figure 12.38 g-AlF3

(a) [111] projection of one HTB layer and (b) [101] projection of the structure of

of the pseudo-hexagonal cell (Figure 12.39b). The structure is also described from perovskite-type layers in (110) (1–10) or (010) planes connected by infinite [001] trans-chains of octahedra (Figure 12.39a). a

b

b

b

a

c

Figure 12.39

12.6.3.3

(a) [100] and (b) [001] projections of the structure of -AlF3

t-AlF3 (or y-AlF3)

Thermal decomposition of [N(CH3)4] (AlF4) H2O[118] under vacuum at 450 C gives a tetragonal variety t-AlF3 (or y-AlF3) with a small amount of -AlF3. The structure, determined from X-ray powder diffraction data, results from the stacking of chiolite type Al3F14 layers (A) (Figure 12.40a) and Al5F24 layers (B, B0 ) (Figure 12.40d) along [001]. The stacking order is A-B-B0 (the superscript 0 indicates a translation or an inversion from B). It is observed that Al4F18 tetramers (Figure 12.40b) are associated with trans-AlF5 chains (Figure 12.40c) in this variety of AlF3. 12.6.3.4

k-AlF3

Thermal decomposition of metastable -NH4AlF4 gives the TTB k-AlF3[111]. TTB type AlF4 layers connect along the c axis in order to leave cavities with triangular, square or pentagonal sections (Figure 12.41).

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates a

b

371

d

c b

b

a

a

Figure 12.40 [001] Projection (a) of the A layer of chiolite, (b) Al4F18 tetramer, (c) trans-chain AlF5 and (d) [001] projection of the Al5F24 layer of t-AlF3

b a

Figure 12.41

12.6.3.5

[001] Projection of the structure of k-AlF3

-AlF3

-AlF3 is the thermodynamically stable form of AlF3. The AlF6 octahedra connect along the a, b, c directions of a cubic ReO3[119]-type cell. This ideal structure is adopted above 725 K while a rhombohedral distortion appears below this temperature (Figure 12.42).

b

a

Figure 12.42

c

Projection of the structure of -AlF3 along [241]

372

12.7

Functionalized Inorganic Fluorides

Evolution of the Condensation of Inorganic Polyanions

As demonstrated previously in the Al(OH)3-guaz-HFaq.-EtOH system (Figure 12.1), the composition of the crystalline precipitates evolves consistently with the concentrations of the starting materials. Three other systems illustrate this feature: the tren system for three different [Al3þ] concentrations, 0.1 mol.L1, 0.5 mol.L1, 1.0 mol.L1, at T ¼ 190 C (Figure 12.43), the guanidinium chloride HguaCl system ([Al3þ] ¼ 0.4 mol.L1, T ¼ 190 C) (Figure 12.44), the dien system ([Al3þ] ¼ 0.5 mol.L1, T ¼ 190 C) (Figure 12.45). It is also well known that the thermal decomposition of hydrates proceeds towards condensed phases, frequently in several steps. This point applies to hybrid organicinorganic aluminofluorides.

12.7.1

Influence of Amine and Aluminum Concentrations

At constant concentrations of Al(OH)3, 3D pyrochlore is always obtained at small HF and amine contents; this feature is consistent with the hydroxylation and the inorganic nature of this pyrochlore: Al2[F0.5(OH)0.5]6 H2O. With tren amine, the crystallization domain of pyrochlore is shifted towards higher Al(OH)3/HF ratios when [Al3þ] increases though the [HF þ F] concentration increases in the solution. It is clear that all solids with isolated (AlF6)3 anions crystallize from solutions with high amine concentrations. It is also clear that the increase of the HF/amine ratio (Figure 12.43 and Figure 12.44) induces the condensation of the AlF6 units. 0D polyanions, extended 1D chains or extended 2D layers appear. With tren amine, condensation occurs by the elimination of one fluorine atom (vertex sharing octahedra in Al2F11, vAl4F18, Al7F30. . .) before the elimination of two fluorine atoms at high HF concentrations (edge-sharing octahedra in Al2F10). This evolution is paralleled by the variation of the amine protonation states, [H3tren]3þ or [H4tren]4þ. Both domains [H3tren]3þ and [H4tren]4þ are separated by dotted lines in Figure 12.43. Two phases with mixed edge þ vertex sharing or mixed protonation states are found at the borders of the preceding domains (Figure 12.43): [H3tren]2 (Al4F18) and [H3tren]2 (Al4F18) 3.5H2O, respectively. The highest condensation is observed in [H4tren]2 (Al7F29) 2H2O where Al7F30 units share two opposite corners in order to build infinite Al7F29 chains. It must be noted that one more phase was once evidenced in the Al2O3-tren-HFaq.-EtOH system, [H4tren]3/2 (Al6F24) 2H2O, but could not be reproduced. In this phase, infinite Al6F24 chains can be visualized from Al7F30 units that share AlF6 octahedra; such chains were previously described in (H3O) [H2pipz] (Al3F12) and g-CsAlF4. Fluoride crystallization is not strongly influenced by the chemical state of the amine. For example, guanidinium chloride can be used instead of guanidine[77]: at [HguaCl] ¼ 1.6 mol.L1 and [Al3þ] ¼ 0.4 mol.L1 (Figure 12.44), the increase of HF concentration implies that [Hgua]3 (AlF6) appears before (H3O)2 [Hgua]16 (Al4F18)3. At lower [HguaCl] concentration (0.17 mol.L1), [Hgua]2 (Al5F17) appears before (H3O)2 [Hgua]16 (Al4F18)3. At constant HF concentration, a decrease of the amine content favours the condensation of the metal species.

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

373

50 tren 50 [Al3+] = 0.1 mol.L–1 60

40

[H3tren]·(AlF6)·HF 70 [H3tren]·(AlF6)·H2O [H4tren]·(AlF6)·(F) 80

30

20 [H3tren]4·(Al8F35)·(OH)·H2O) 10 Al2(F0.5,OH0.5)6·H2O

90 (H3O)·H4tren]2·(Al7F30) [H4tren]·(Al2F10)·H2O 100 [H tren]2·(Al7F29)·2H2O HF 0 4 10 20

0 50 Al(OH)

30

40 tren 50 50

[Al3+] = 0.5 mol.L–1 60 [H3tren]·(AlF6)·HF [H3tren]·(AlF6)·H2O [H4tren]·(AlF6)·(F) 70 (H3O)·[H4tren]2·(AlF6)3·6H2O [H3tren]2·(AlF5(H2O))3·8H2O

3

40

30

80

20

[H3tren]4·(Al2F11)·(AlF6)2·(F)·10H2O 90

[H3tren]4·(Al8F35)·(OH)·H2O 10 Al2(F0.5,OH0.5)6·H2O

[H4tren]·(Al2F10)·H2O 100 HF 0 [Al3+] = 1 mol.L–1

0 50 Al(OH)

[H4tren]2·(Al7F29)·2H2O 10 20 30 40 50 tren 50

[H3tren]·(AlF6)·HF [H3tren]·(AlF6)·H2O [H4tren]·(AlF6)·(F) 70 (H3O)·[H4tren]2·(AlF6)3·6H2O [H3tren]2·(AlF5(H2O))3·8H2O 80

90 [H4tren]·(Al2F10)·H2O

60

3

40

30

20 [H3tren]4·(Al8F35)·(OH)·H2O 10 Al2(F0.5,OH0.5)6·H2O

100 HF 0

[H4tren]2·(Al7F29)·2H2O 10 20 30 [H3tren]2·(Al4F18)·3.5H2O [H3tren]2·(Al4F18)

40

0 50 Al(OH)

3

Figure 12.43 Composition space representations of the Al(OH)3-tren-HFaq.-EtOH system for [Al3þ] ¼ 0.1 mol.L1, [Al3þ] ¼ 0.5 mol.L1, [Al3þ] ¼ 1.0 mol.L1 and T ¼ 190 C

374

Functionalized Inorganic Fluorides 50

HguaCl 50

[Al3+] = 0.4 mol.L–1 60

40

70

30

[Hgua]3·(AlF6)

Al(OH)3 20

80

90

Al2(F0.5,OH0.5)6·H2O 0 40 50 Al(OH)3

100 HF

0

10

[Hgua]2·(Al5F17)

(H3O)2·[Hgua]16·(Al4F18)3

10

20

30

Figure 12.44 Composition space representation of the Al(OH)3-HguaCl-HFaq.-EtOH system for [Al3þ] ¼ 0.4 mol.L1 and T ¼ 190 C

[Al3+] = 0.5

mol.L–1

60 70 [H3dien]·(AlF6) 80 [H3dien]·(AlF6)·2H2O 90 100 0

40 30 Al(OH)3 20 10

[H3dien]2·(AlF5(H2O))3·2H2O HF

50 dien 50

10

Al2(F0.5,OH0.5)6·H2O 0 [H3dien]2·(Al4F18) 20 30 40 50 Al(OH)3

Figure 12.45 Composition space representation of the Al(OH)3-dien-HFaq.-EtOH system for [Al3þ] ¼ 0.5 mol.L1 and T ¼ 190 C

These general trends are helpful to establish the opportunity of a difficult ab initio structure determination from powder diffraction data.

12.7.2

Temperature

The evaporation of a solution of Al(OH)3, guaz, HFaq., H2O in 1/5/4.5/55.5 ratio at room temperature leaves transparent crystals of the hydrate [Hguaz]2 (AlF5(H2O)) 2H2O. On

Polyanion Condensation in Inorganic and Hybrid Fluoroaluminates

375

heating, this hydrate decomposes into several steps to give, as proved by X-ray thermodiffraction (Figure 12.46), [Hguaz]2 (AlF5) at T 100–150 C, [Hguaz]2 (Al2F8) at T 210 C and, finally, -AlF3 at T 290 C. Then, the transformation of this hydrate to the 1D [Hguaz]2 (AlF5) and 1D [Hguaz]2 (Al2F8) follows exactly the evolution of the crystalline phases with the composition of the solutions (Figure 12.47): the thermal condensation path is parallel to the chemical reaction path in the composition space diagram [41].

α-AlF3 290°C

[Hguaz]2·(Al2F8) 210°C [Hguaz]2·(AlF5) 150°C [Hguaz]2·(AlF5(H2O))·2H2O 25°C 5

10

Figure 12.46

20

30

X-ray thermodiffraction of [Hguaz]2 (AlF5(H2O)) 2H2O

[Al3+] = 1 mol.L–1

guaz 50

50 [Hguaz]2·(AlF5(H2O))·2H2O

60

[Hguaz]2·(AlF5)

70 80

40 30

[Hguaz] [H 2·(Al2F8)

90 [Hguaz]2·(Al5F17)] 100 HF 0

10

20

20 10 AlAl 2(Fe,OH0.5)6·H2O

30

40

0 50 Al(OH)

Figure 12.47 Composition space representation of the Al(OH)3-guaz-HFaq.-EtOH system for [Al3þ] ¼ 1 mol.L1 and T ¼ 190 C

376

Functionalized Inorganic Fluorides

Acknowledgements The authors wish to thank A. Ben Ali, A. Jouanneaux, J. Lhoste and M.A. Saada.

Supplementary Materials A data file of all hybrid and inorganic fluoroaluminates can be freely downloaded on the website (Te´le´chargements): www.univ-lemans.fr/sciences/fluorures/ldf.html.

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13 Synthesis, Structure and Superconducting/Magnetic Properties of Cu- and Mn-based Oxyfluorides Evgeny V. Antipov and Artem M. Abakumov Department of Chemistry, Moscow State University, Moscow, Russia

13.1

Introduction

Complex transition metal oxides demonstrate a large variety of practically important physical properties and, at the same time, provide an excellent opportunity to investigate the fundamental relationships between crystal structure, electronic state of the transition metal cations, magnetic and electric transport properties. In spite of the different physics behind such phenomena as high-temperature superconductivity, magnetoresistance, magnetoelectricity etc., they are all related to strongly correlated electronic systems. Such systems are usually very structure-sensitive and their behaviour is governed by the parameters determining the exchange interactions between the transition metal cations (geometry and anisotropy of a metal-anion-metal bonding framework, metal-anion bond distances and angles), electronic state of the transition metal cations (often expressed in terms of formal oxidation state) and local structure distortions (breaking local symmetry of coordination environment of the transition metal cation, caused, for example, by a disordered placement of different species at the cation or/and anion sublattices). Target-aimed synthesis of the complex transition metal oxides and their derivatives requires taking their great complexity into account, this being caused by variations in the Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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valent state of transition metal cations due to heterovalent replacements or vacancies formation at the cation or anion sublattices, ordering of the different species in both sublattices, structural distortions due to electronic instabilities (static Jahn-Teller effect), charge and orbital ordering, and defects (stacking faults, translational and orientational domains). This chapter is focused on the structural peculiarities of fluorine insertion into the anion sublattice of complex cuprates and manganites, in relation to their magnetic and electric transport properties (particularly, superconductivity for the layered cuprates). Indeed, the Cu2þ and Mn3þ cations demonstrate a certain similarity in their crystal chemistry. Both are Jahn-Teller active cations due to asymmetric occupation of the eg orbitals (t2g6eg3 for Cu2þ and t2g3eg1 for Mn3þ), which results in a distortion of their octahedral coordination environment. The ionic radii of Cu2þ and Mn3þ in a high-spin ˚ , r(Cu2þ) ¼ 0.87 A ˚ ) [1]. For Cu2þ, as well as for state are also close (r(Mn3þ) ¼ 0.79 A 3þ Mn , there exists a tendency to remove some of the anions from the octahedral coordination environment resulting in tetragonal pyramidal or square planar coordination (the latter is more characteristic of Cu2þ than of Mn3þ). Due to this similarity, some of the perovskite-based manganites have their counterparts among layered cuprates. Famous high-temperature superconductors (HTSC) La2xAxCuO4 (A ¼ Sr, Ba) belong to the Ruddlesden-Popper homologous series Anþ1BnO3nþ1 (n ¼ 1), as well as LaSrMnO4 manganite [2,3]. Using the analogy with the HTSC cuprates, the layered manganites BiPbSr2MnO6þ, Bi2Sr2MnO6þ (2201 phases) [4–6] and Bi2xPbxSr1.5Ca1.5Mn2O9 (2212 phase) [7] were prepared. One can expect that certain similarity will be also observed with respect to formation of oxyfluorides. Crystal chemistry of oxygen and fluorine anions is also similar; they both have high ˚ , rO2- ¼ 1.35 A ˚ ) [1] and oxides and fluorelectronegativity, close ionic radii (rF- ¼ 1.29 A ides often have the same structures (CaO and NaF, LaOF and CaF2, La2CuO4 and K2CuF4 etc.). Hence, the formation of oxyfluorides with the crystal structures close to those of corresponding oxides is possible and this possibility can be used for adjusting the properties of oxides.

13.2

Chemical Aspects of Fluorination of Complex Oxides

Soon after the discovery of high-temperature superconductivity in layered cuprates [8], the first attempts were undertaken to vary the formal copper valence (VCu) and optimize the superconducting properties by insertion of fluorine into the cuprate structures. These attempts were based on a solid state reaction, where fluorine was introduced into an initial mixture of oxides and/or carbonates by means of a partial replacement of one of the reagents by corresponding alkali-earth metal fluoride, rare-earth fluoride or oxyfluoride or CuF2. Using this technique with NdF3 [9–13] as a source of fluorine, the first electrondoped superconductor (VCu < þ2) Nd2CuO4xFx was prepared at 860–1150 C. Hole-doped superconductors (VCu > þ2) Sr2RCu2O5F1þ (R ¼ Y, La, Nd, Sm, Gd, Dy, Er, Yb) [14], Sr2CaCu2O4þF2–y and Sr2Ca2Cu3O6þF2–y [15, 16] were obtained through a high-temperature high-pressure solid-state reaction (5.5–6 GPa, 1250–1300 C), where fluorine was introduced in a form of SrF2 or CaF2. The high-pressure high-temperature

Properties of Cu- and Mn-based Oxyfluorides

385

technique (3.5–4.5 GPa, 900–1050 C) was also successful in the preparation of the cuprates belonging to the Ba2Can1CunO2n(O,F)2 (n ¼ 2–5) homologues series [17–19]. However, at ambient pressure the high-temperature solid-state reaction does not lead to fluorinated derivatives of the hole-doped complex cuprates, especially if the compound contains alkali earth metal cations. For example, different reports on an increase of Tc above 100 K for fluorinated YBa2Cu3OyFx (Y123), obtained by a solid-state reaction with BaF2 have never been reproduced. Moreover, thorough investigation of the products of solid-state reaction aimed at fluorinated Y123 demonstrated that BaF2 is always present in the reaction mixture obtained by high temperature annealing at 800–950 C, independently from the form in which fluorine was present in the initial reagents [20–23]. High-temperature solid-state reaction cannot be used for the preparation of other HTSC cuprates, such as Sr2CuO2F2, which decomposes at temperatures 480 C producing SrF2, Sr2CuO3 and Sr14Cu24O41 [24,25]. However, Sr2CuO2F2þ can be prepared under much ‘softer’ conditions by low-temperature treatment of the parent Sr2CuO3 phase by various fluorinating agents: a mixture of 10 %F2 and 90 %N2 (15 min, t ¼ 210 C) [26], NH4F in a presence of O2 (6–8 h, t ¼ 225 oC) [25, 27, 28] or XeF2 (40 min – 10 h, t ¼ 160 –350 C) [24]. The impossibility of using a high-temperature solid-state reaction for preparation of fluorinated cuprates can be related to high lattice energies of their decomposition products, namely fluorite-like fluorides of alkali-earth metals AF2 and rare-earth oxyfluorides ROF. This can be reflected by high enthalpy of the reaction between Y123 and gaseous fluorine: YBa2 Cu3 O6 þ 5=2 F2 ! YOF þ 2BaF2 þ 3CuO þ O2 DHo ¼1518 kJ Various fluorinating agents with either an oxidizing or a reducing character have been used for low-temperature synthesis of complex copper oxyfluorides: gaseous fluorine (pure [29–31] or diluted with other gases [26, 32–34]), NF3 [35, 36], NH4F [31, 37–40], NH4HF2 [42–44], transition metal fluorides (ZnF2, CuF2, NiF2, AgF2 [31, 45–49]), ClF3 (here fluorine is incorporated along with chlorine) [50–53]. NH4F and NH4HF2 are nonoxidizing fluorinating agents, although they allow superconductors with formal copper oxidation state >þ2 to be obtained in the presence of oxygen due to reactions [28]: ð2 þ ÞNH4 F þ ð3 þ 2Þ=2O2 þ Sr2 CuO3 ! Sr2 CuO2 F2 þ þ ð2 þ Þ=2N2 þ2ð2 þ ÞH2 O ð2 þ ÞNH4 F þ =4O2 þ Sr2 CuO3 ! Sr2 CuO2 F2 þ þ ð2 þ Þ=2H2 O þ ð2 þ ÞNH3 : Moreover, a mechanism was proposed that allows NH4F to be an oxidizing fluorinating agent even in an oxygen-free atmosphere. The reaction of the parent complex oxide with NH4F produces H2O as by-product. Released water then decomposes the obtained fluorinated oxide due to pyrohydrolysis. It is assumed that pyrohydrolysis leads to amorphous compounds containing mostly Cuþ, so that the increase of formal copper valence in the superconducting oxyfluoride is explained by disproportionation: Cu(II) (initial oxide) ! Cu(I) (pyrohydrolysis products) þ Cu(II þ III) (superconducting oxyfluoride) [27]. The target material is significantly contaminated by pyrohydrolysis products. Fluorides of metals, which do not form thermally stable oxides, such as AgF2, can serve as oxidizing fluorinating

386

Functionalized Inorganic Fluorides

agents, but along with fluorine insertion into the initial oxide structure, oxygen can be partially replaced by fluorine, especially at the elevated temperature of the reaction. Fluorination by means of transition metal fluorides MF2 (M ¼ Zn, Cu, Ni) occurs according to the reaction: MF2 þ Bm A2 A 0n 1 Cun O2n þ 2 þ mx ! MO þ Bm A2 A 0n 1 Cun O2n þ 2 þ mx F2 ; Here the general CmA2A0 n1CunO2nþ2þmx formula of layered cuprates is used according to the layer sequence – (AO) – (BOx)m – (AO) – (CuO2) – (A0 CuO2)n1 – (AO) – (BOx)m – (AO) –, where 0 £ x £ 1, m ¼ 0, 1, 2, n ¼ 1, 2, 3. . . A, A0 – alkali-earth or rare-earth cations, B – smaller cation, such as Hg2þ, Bi3þ. This reaction implies that copper oxidation state is unchanged. If fluorination does not occur via gas phase, i.e. if MF2 is mixed with the initial oxide, the final product will be contaminated with the metal oxide MO. XeF2, proposed in [24] for fluorination of complex cuprates, has several advantages in comparison with other fluorinating agents. XeF2 is an oxidizing fluorinating agent. It is solid at a room temperature, which allows its quantity to be measured according to the required degree of fluorination and oxidation state of the transition metal. It is less aggressive than gaseous F2 and does not require a complicated experimental setup for fluorination. Besides the compounds resulting from interaction between XeF2 and the initial oxide, no other side compounds are produced, preventing additional admixtures in the sample. The reaction with XeF2 should be performed in water-free conditions. Even traces of water decrease drastically the purity of the final product because of pyrohidrolisis. HF, releasing from interaction of XeF2 and obtained copper oxyfluoride with H2O, further decomposes the initial oxide and oxyfluoride producing water, which can again participate in pyrohydrolisis reaction [24]: Sr2 CuðO; F Þ 4 þ þ H2 O ! SrF2 þ Sr2 CuO3 þ CuO þ HF Sr2 CuO3 ðor Sr2 CuðO;F Þ 4 þ Þ þ HF ! SrF2 þ CuO þ H2 O Due to sensitivity of the reaction to water traces and high volatility of XeF2 at elevated temperatures, fluorination should be done in hermetically closed vessels, which are stable with respect to fluorine affection. It can be done in a sealed Ar-filled Cu tube containing Ni crucible with a mixture of the initial oxide and XeF2. All preparation should be performed in a glove box with a water-free atmosphere. The fluorination reaction with XeF2 occurs in a temperature interval of 150–400 C. Under these conditions, XeF2 is in a gaseous phase and interaction with the material occurs through the solid-gas interface. According to the temperature dependence of equilibrium composition of the gas mixture at 1 atm. containing Xe and F2 in 1:1 molar ratio [54], main constituents at temperature < 400 C are Xe, XeF2 and XeFe4 (Figure 13.1). Both XeF2 and XeF4 can act as fluorinating agents. The oxidizing potential of XeF4 is higher than that of XeF2 [55,56], but its concentration is much lower. Above 400 C dissociation into Xe and F2 becomes noticeable. Interaction between complex oxide and XeF2 can occur according to different reaction schemes, depending on the temperature of the reaction. In [57] it was demonstrated that fluorination of the YBa2Cu3O6.11 phase at temperatures between 150 C and 300 C occurs as a fluorine insertion: YBa2 Cu3 O6:11 þ 1=2x XeF2 ! YBa2 Cu3 O6 þ Fx þ 1=2x Xe

Properties of Cu- and Mn-based Oxyfluorides

387

1.0

0.8

Molar fraction

XeF2 0.6

0.4 Xe F2

0.2

XeF4 0.0 200

300

400 t, °C

500

600

Figure 13.1 Temperature dependence of equilibrium composition of the gas mixture containing 1 mol of XeF2 at 1 atm. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

As a result, the formal copper oxidation state increases from VCu ¼ þ1.74 in the initial strongly reduced YBa2Cu3O6.11 phase to VCu ¼ þ2.03 – þ2.14 in the fluorinated samples. Increasing the temperature to 400 C results in a competing anion exchange reaction, that is a replacement of part of oxygen atoms with fluorine. Anion exchange causes partial reduction of copper cations because one N2 anion in the initial oxide is replaced by one F in the fluorinated phase. It is difficult to find experimental conditions for either exclusively fluorine insertion or anion exchange reactions. Usually these two competing processes occur simultaneously: YBa2 Cu3 O6 þ 1=2 ðy þ xÞ XeF2 ! YBa2 Cu3 O6 y Fx Fy þ 1=2 ðy þ xÞ Xe þ y½O Increasing temperature favours the anion exchange. Interaction of the oxidized YBa2Cu3O6.95 phase with XeF2 at 350 C decreases the copper oxidation state from VCu ¼ þ 2.30 to VCu ¼ þ 2.22. Oxygen [O], releasing in the anion exchange reaction is absorbed by the inner walls of the copper tube serving as a reaction vessel, forming Cu2O. This oxygen cannot act as an oxidizer because its partial pressure in the tube (corresponding to an equilibrium partial oxygen pressure under Cu/Cu2O mixture at the reaction temperature, lnP(O2) ¼ –22.2 at R ¼ 600 K) is comparable with the partial oxygen pressure under strongly reduced YBa2Cu3O6.1 phase at the same temperature (lnP(O2) ¼ 21.1) [57]. Because different temperature intervals are preferable for the fluorine insertion and anion exchange reactions, it can be used for performing fluorination in either oxidizing or reducing regimes, which can lead to either p- or n-type doped superconductors. However, a large

388

Functionalized Inorganic Fluorides

degree of anion exchange should be avoided because the final stage of this reaction, occurring at 400–500 C is a decomposition of the complex cuprate: YBa2 Cu3 O6 þ 5=2 XeF2 ! YOF þ 2BaF2 þ 3CuO þ 5=2 Xe þ 2½O YBa2 Cu3 O6 þ 5=2 F2 ! YOF þ 2BaF2 þ 3CuO þ 2½O Recently, a poly(vinylidene) fluoride (PVDF) was proposed as an efficient fluorinated agent. Being tested on a preparation of Ca2CuO2F2 by a fluorination of Ca2CuO3 at 350 C, it demonstrated superior performance in comparison with NH4F or F2 [58]. It is assumed that volatile decomposition products of PVDF, such as COF2, interact with the material through the gas phase, producing the oxyfluoride [59]. Basically the same synthesis pathways were applied for the preparation of the fluorinated derivatives of complex manganites. XeF2 was used for fluorination of the Sr2Mn2O5 anion-deficient perovskite [60] and Mn-based brownmilerites [61,62], whereas the fluorine-containing derivatives of the Ruddlesden-Popper manganites were prepared with 10 %F2 þ 90 %N2 gas mixture [63,64], CuF2 [37] or PVDF [59]. In all cases fluorination was performed at a moderately low temperature, preserving the cation arrangement of the structure of the initial oxide. An interesting synthesis approach should be mentioned, which implies a preparation of highly fluorinated LaSrMnO4F1.7 or La1.2Sr1.8Mn2O7F2 phases and then a solid state reaction between the fluorinated and the initial material at 300–500 C to produce a compound with intermediate desired amount of fluorine, such as LaSrMnO4F or La1.2Sr1.8Mn2O7F [63, 64].

13.3

Structural Aspects of Fluorination of Complex Cuprates and Superconducting Properties

Investigations of the relationships between the composition, crystal structure and superconducting properties of different families of layered cuprates resulted in empirical criteria that have to be fulfilled for superconductivity to appear [65]: • layered structure with flat or slightly deformed conducting (CuO2) planes consisting of corner-sharing CuO4 squares; ˚ for efficient • optimal interatomic Cu-O distances within the (CuO2) plane of 1.90–1.97 A overlap of the 3dx2y2 copper orbitals and 2px,y oxygen orbitals resulting in delocalized states in the * band; • optimal concentration of charge carriers (electrons or holes) in the (CuO2) planes, corresponding to the range of formal copper oxidation state from þ2.05 to þ2.25 for the hole-doped superconductors and from þ1.8 to þ1.9 for the electron-doped superconductors. Formal negative net charge of the (CuO2) planes is compensated by neighbouring insulating positively charged blocks, which serve as ‘charge reservoirs’ for variation of the formal Cu oxidation state. The positive charge of these blocks can be changed either by heterovalent cation replacement or by changing oxygen content. An alternative way is a replacement of some of the oxygen atoms by fluorine, due to different formal charges of O2 and F. Fluorine can also be introduced into the structure to occupy vacant

Properties of Cu- and Mn-based Oxyfluorides

389

anion positions. Both cases result in changing formal Cu oxidation states and, if the (CuO2) planes are not significantly disturbed, the oxyfluoride can demonstrate superconducting properties. Formally, different modes of tuning charge carrier concentration can be distinguished. Fluorine can be inserted into the vacant anion position of complex cuprates: Bm A2 A 0n 1 Cun O2n þ 2 þ mx þ ½F ! Bm A2 A 0n 1 Cun O2n þ 2 þ mx F that results in a hole doping. Fluorine can also substitute equal amount of oxygen in the structure: Bm A2 A 0n 1 Cun O2n þ 2 þ mx þ ½F ! Bm A2 A 0n 1 Cun O2n þ 2 þ mx F þ ½O that leads to electron doping. Both processes can occur simultaneously resulting in a replacement of one oxygen atom by two fluorine ones: Bm A2 A 0n 1 Cun O2n þ 2 þ mx þ ð þ 0 Þ½F ! Bm A2 A 0n 1 Cun O2n þ 2 þ mx Fþ 0 þ ½O that does not alter the copper valence, but increases amount of anions in the structure. Besides tuning charge carrier concentration, fluorination provides extra possibilities for searching new layered compounds. The simultaneous requirements of preserving electroneutrality and geometrical matching between the (CuO2) planes and the ‘charge reservoir’ blocks impose severe restrictions on possible types of cations used. Decreasing a formal charge of the anion sublattice by the O2 ! F replacement can expand the range of cations for a formation of layered structure. At the same time a twice smaller formal charge of fluorine anion requires a twice larger amount of overstoichiometric anions for the same level of hole doping. Increasing the amount of anions itself can cause significant structure rearrangements followed by changing superconducting properties.

13.3.1

Electron Doped Superconductors: Heterovalent Replacement 1O2 ! 1F

Nd2CuO4 has a tetragonal T’ structure, which can be represented as an ordered alternation of flat (CuO2) planes and (Nd2O2) blocks with a fluorite-type structure, where the oxygen atoms occupy tetrahedral interstices in the close packed arrangement of the Nd atoms [66] (Figure 13.2). The in-plane Cu-Oeq distances in this structure are relatively long ˚ ), which is suitable for superconductivity due to electron doping. Reduction of (1.972 A the copper valence below þ2 was achieved, for example, by Nd3þ ! Ce4þ heterovalent replacement (Nd2xCexCuO4, Rc ¼ 25 K at 0.14 < x < 0.18) [67, 68]. Heterovalent replacement of 1O2 by 1F leads to a similar result. The Nd2CuO4xFx solid solution exists up to x 0.3 [10] and Tc reaches its maximum of 27 K in the range of 0.2 £ x £ 0.3 [9, 10, 38]. According to electrochemical analysis of copper oxidation state, the Nd2CuO3.74F0.26 composition corresponds to Tc ¼ 27 K [69]. Fluorination of Nd2CuO4 does not lead to drastic structural changes. It is accompanied by the enlargement of the a unit cell parameter because reduction of the copper atoms is related to increasing occupation of the anti-bonding *-band and increasing length of the Cu-Oeq bond. The structure slightly shrinks along the c-axis because of smaller ionic radius of F in

390

Functionalized Inorganic Fluorides Cu CuO2

Nd2O2 Nd CuO2

Nd2O2

CuO2

Figure 13.2 The crystal structure of Nd2CuO4. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

comparison to that of O2 [10,38]. X-ray absorption fine structure (XAFS) studies on the Nd2CuO4xFx detected the majority of F atoms to substitute oxygens at the tetrahedrally coordinated positions of the fluorite-type (Nd2O2) block. Although up to 6 % of oxygen anions were also substituted in the (CuO2) planes, it did not destroy the superconducting properties [70]. Preferential location of fluorine in the tetrahedrally coordinated sites was also supported by 19F NMR spectroscopy [71]. Superstructures in Nd2CuO4, Nd2CuO4xFx and Nd2xCexCuO4 were investigated using electron diffraction (ED) and transmission electron microscopy (TEM) [72]. The observed superlattices were identical for all three materials, which reflects that these superlattices are not related to a replacement of oxygen by fluorine. TEM study of the Nd2CuO3.8F0.2 sample demonstrated absence of superstructure for this compound [10]. Pr2CuO4 under fluorination behaves similar to its Nd-based analogue, reaching Tc ¼ 27–30 K [73, 74].

13.3.2

Hole Doped Superconductors: Fluorine Insertion into Vacant Anion Sites

Vacant positions available for oxygen cause wide-range oxygen non-stoichiometry for many HTSC cuprates related to changes in charge carrier concentration and superconducting properties. Such positions are also available for fluorine. According to a simple ionic model of doping, twice larger amounts of the F anions are required compared to the O2 anions for the same doping level. Increasing occupation of the anion positions can be the origin of substantial structural differences between the compounds doped with either oxygen or fluorine. Comparison of the crystal structures of the fluorinated and oxygenated materials often provides valuable information on the doping mechanisms in the cuprates.

Properties of Cu- and Mn-based Oxyfluorides

391

Perovskite-like anion deficient structure of the RBa2Cu3O6þ phases (R123, R – rareearth cation) (a b ap,c 3ap, ap – the parameter of the perovskite subcell) consists of the sequence of the layers along the c-axis: ðCuNO Þ ðBaOÞ ðCuO2 Þ ðRNÞ ðCuO2 Þ ðBaOÞ ðCuN O Þ ðNanion vacancyÞ: All anion positions in the (CuNO) layer (also called the Cu1 layer) can be vacant ( ¼ 0) in the tetragonal structure of strongly reduced RBa2Cu3O6 (Figure 13.3), which does not demonstrate superconducting properties (VCu ¼ þ1.67). The Cu1 atoms are in dumbbell coordination, suitable for the Cuþ cations. At 1 oxygen anions occupy half of the positions in the Cu1 layers in an ordered manner, thus completing the coordination environment of the Cu1 cations to a square (Figure 13.3). Chains of corner-sharing CuO4 squares run along the b-axis resulting in an orthorhombic distortion. Fully oxygenated compounds with 1 have the optimal hole concentration for this type of structure and demonstrate a highest Tc of 93–94 K. YBa2Cu3O7

YBa2Cu3O6

-Ba

Figure 13.3

-Y

-Cu

YBa2Cu3O6F2

-O

-F

The crystal structures of YBa2Cu3O7, YBa2Cu3O6 and YBa2Cu3O6F2

Fluorine enters into the structures of partially reduced or fully oxygenated YBa2Cu3O6þ ( 0.65) occupying vacant anion positions in the Cu1 layers along both a and b axes of orthorhombic unit cell. At a low degree of fluorination, fluorine occupies the (0,1/2,0) position preferentially but when the occupancy factor of this position reaches 80 %, it enters into the position (1/2,0,0) [75–77]. This order of filling the vacant positions agrees with an increase of the orthorhombic distortion in YBa2Cu3O6.7F0.15 and its subsequent decrease in YBa2Cu3O6.7F0.3. Due to similar environment of the (0,1/2,0) and (1/2,0,0) positions it is not possible to distinguish fluorine distribution using 19F NMR spectroscopy [78, 79]. Fluorination of YBa2Cu3O6.7 is accompanied by rising Tc from 65 K to 90 K [80]. Fluorination of fully oxygenated YBa2Cu3O6þ with close to 1 in the majority of cases results in the oxyfluorides with Tc not higher than that for

392

Functionalized Inorganic Fluorides

the sample optimally doped with oxygen [29,35, 36, 81 - 84] or in slight increase in Tc from 92 K to 96 K [85]. Fluorination of the YBa2Cu3O6.94 with ZnF2 occurs as an anion exchange reaction resulting in the YBa2Cu3O6.6F0.4 oxyfluoride, which does not demonstrate superconducting properties, most probably due to high degree of fluorine for oxygen substitution in the (CuO2) planes [86]. Fluorine insertion into the strongly reduced Y123 ( 0) occurs more easily than replacement of oxygen by fluorine in the fully oxidized Y123 ( 1) [87]. Fluorination of the strongly reduced tetragonal YBa2Cu3O6.2 leads to an orthorhombic distortion and appearance of a broad superconducting transition with Rc ¼ 55–60 K, which is close to Tc for YBa2Cu3O6.5 [81]. An increasing amount of fluorine is accompanied by an orthorhombic to tetragonal transition. The c parameter of the tetragonal phase ˚ ) is significantly smaller that that for the initial compound (c ¼ 11.839 A ˚ ). (c ¼ 11.640 A This is related to a relatively high residual concentration of anion vacancies in the Cu1 layer and random anion distribution among the (1/2,0,0) and (0,1/2,0) positions. Independently from the initial oxygen content in the R123 phases, increasing the fluorination degree decreases the orthorhombic distortion and finally leads to the tetragonal oxyfluoride [81,88]. Fluorination is accompanied by shrinkage of the structure along the c-axis, similar to the effect of oxygen insertion. It confirms that the coordination environment of the Cu1 atoms is virtually identical in the fluorinated and oxygenated samples and significant amount of anion vacancies remain in the Cu1 layers. Tetragonal symmetry arises from random occupation of the (1/2,0,0) and (0,1/2,0) positions by oxygen and fluorine (n ¼ 0.68), as was found for YBa2Cu3O6.65, treated with NF3 [88]. Starting from simple crystal chemistry considerations Kistenmacher [89] predicted existence of the complex oxyfluoride YBa2Cu3O6F2, where all anion vacancies in the Cu1 layers are filled with fluorine atoms. The copper oxidation state in this compound should be þ2.33, which corresponds to VCu in the oxygen-doped compound YBa2Cu3O7 with highest Tc. This oxyfluoride should possess tetragonal symmetry because the Cu-O chains running along the b-axis are replaced by two-dimensional (CuF2) layers. It could be expected that the c unit cell parameter of YBa2Cu3O6F2 should be significantly larger than that for the oxygenated phase because of the apical elongation of the CuO2F4 octahedra due to Jahn-Teller effect intrinsic in the Cu2þ cations. In order to obtain the YBa2Cu3O6F2 oxyfluoride the initial compound should contain as less oxygen as possible in the Cu1 layers. Fluorination of the non-superconducting YBa2Cu3O6.11 with XeF2 resulted in an increase of VCu to þ2.14 and appearance of superconductivity with Tc ¼ 94 K with a relatively large fraction of the superconducting phase (25 %). Powder XRD pattern of this material demonstrates narrow hk0 reflections of the tetragonal unit cell, whereas the reflections with l 6¼ 0 are practically absent. This implies a strong disorder along the c-axis [57]. A similar effect (suppression of periodicity along the c-axis, corresponding to tripled perovskite unit cell) was also observed in the YBa2Cu3O6.93x/2Fx solid solutions at x ¼ 1.95 [90]. Electron diffraction and high-resolution transmission electron microscopy (HRTEM) images revealed that a new tetragonal ˚ , c 13 A ˚ was obtained [91]. Domains fluorinated phase with lattice parameters a 3.87 A of this highly fluorinated F2 phase alternate with the domains of the initial structure with a smaller amount of fluorine (F1 phase). The meaning is that the domains are of large size in the ab plane and have limited extension along the c-axis that causes anisotropic broadening of the reflections (particularly for the reflections with l 6¼ 0, whereas the hk0 reflections

Properties of Cu- and Mn-based Oxyfluorides

393

stay sharp) on the powder XRD pattern. A significant increase in the c-parameter of the F2 ˚ to 13 A ˚ ) is related to phase in comparison with the initial compound (from 11.815 A increasing distance between the BaO and Cu1 layers. It is in agreement with fluorine insertion into all vacant sites of the Cu1 layers and the formation of apically elongated CuO2F4 octahedra (Figure 13.3). The observed expansion of the structure along the c-axis ˚ in YBa2Cu3O6.11 to results from increasing the apical Cu-Oap distance from 1.85 A ˚ 2.3–2.5 A in the F2 phase with tentative YBa2Cu3O6F2 composition. Another high-Tc superconductor with a structure closely related to that of the Y123 compound is Y2Ba4Cu7O14þ (Y247 phase). Its crystal structure can be described as an ordered alternation of blocks comprising single CuO chains (Y123-units) and double chains of the edge sharing CuO4 squares (Y124 or YBa2Cu4O8 units) along the c-axis (Figure 13.4). The Y247 phase exhibits a superconducting transition with Tc ranging from 14 K to 95 K, depending on the oxygen content. Oxygen non-stoichiometry in the Y247 compound is related to the presence of the Y123 blocks. Fluorine insertion into the structure of the strongly reduced Y2Ba4Cu7O14.09 phase results in increasing Tc from ˚ ) is signifi30 K to 62 K [92]. The unit cell c-parameter of the fluorinated phase (53.5 A cantly larger than that of the initial compound. However, in contrast to the fluorinated YBa2Cu3O6F2 phase, the structure of the fluorinated Y247 remains orthorhombically distorted because fluorine insertion does affect the double chains. -Ba -Y -Cu -O

c /2

Figure 13.4

The crystal structure of Y2Ba4Cu7O14þ (=0).

According to the HRTEM results, fluorine insertion into the Y123 and Y247 structures occurs inhomogeneously. The fully fluorinated YBa2Cu3O6F2 phase was observed as domains in the matrix of partially fluorinated material, whereas in fluorinated Y247 the areas with complete occupation of all anion positions in the Cu1 layers by fluorine were observed only as local ‘pancake’-type defects. Similar microstructure was found in fluorinated RBa2Cu3O6.12 (R123, R ¼ Dy, Ho h Tm) phases [93]. In contrast to Y123,

394

Functionalized Inorganic Fluorides

interaction of R123 with XeF2 occurs mainly through anion exchange, which results in a copper oxidation state that is too low for the appearance of superconductivity. The structure of PbYBa2Cu3O7 is close to that of Y123 and can be derived from Y123 structure by insertion of (PbO) layer between the (CuNO) and (BaO) layers. As in Y123, fluorine enters into the (CuNO) layers of the PbYBa2Cu3O7 structure filling all vacant positions and yielding the fluorinated PbYBa2Cu3O7F2 phase, containing layers of the CuOF4 tetragonal pyramids [33]. PbYBa2Cu3O7F2 is superconducting (Tc 30 K), and 20 % Ca-doping on the Y site results in an increase in Tc up to 50 K. Vacant anion positions responsible for wide range oxygen non-stoichiometry exist also in the structures of the Hg-bearing superconducting cuprates belonging to the HgBa2Can1CunO2nþ2þ homologous series. The structures of these compounds can be described by the layer sequence: ðHgO Þ ðBaOÞ ðCuO2 Þ fðCaNÞðCuO2 Þgn1 ðBaOÞ ðHgO Þ where the (BaO)(HgO)(BaO) slabs represent the NaCl-type structure (Figure 13.5). The ˚ . The Hg atoms in the (HgO) layers form a square mesh with a period of 3.9 A overstoichiometric oxygen atoms O are located at the centres of the Hg squares. This position is also available for fluorine atoms. The anions in the (HgO) layers are separated ˚ . Due to stable dumbbell from the Hg cations by relatively large distance of 2.7 A 2þ coordination for the Hg cations and relatively large separation between the Hg atoms and neighbouring anions, the anions are weakly bonded to the Hg atoms. In contrast to Y123 phase, where fluorine atoms significantly change Cu-(O,F) interatomic distances due to formation of strong bonds to the Cu cations, the overstoichiometric anions in the Hg-bearing superconductors are weakly bonded to the Hg cations (Figure 13.6). Increasing the amount of anions in the Hg-layers with twice as many F atoms as O ones does not drastically affect the crystal structure but results in a variation in interatomic distances keeping the same doping level. By comparing extra anion concentration, Tc values and interatomic distances for oxygenated and fluorinated samples valuable informations has been obtained on the doping mechanism in the Hg-bearing superconductors. Hg-1201

Hg-1212

Hg-1223

Ba Hg CuO2

Ca Ca

CuO2

CuO2 Ca CuO2

Hg Ba

Figure 13.5 The crystal structures HgBa2Ca2Cu3O8þ (from left to right)

of

HgBa2CuO4þ,

HgBa2CaCu2O6þ

and

Properties of Cu- and Mn-based Oxyfluorides Hg-1201, Hg-1223

Y-123, Y-247

Cu O

Cu

Cu

Cu Cu

Cu

Cu Cu

F

Cu Cu

395

F

Hg

Cu Cu

Hg

Hg Hg

Figure 13.6 Fluorine insertion schemes into the structures of Y123, Y247 and Hg1201, Hg1223

Attempts were made to introduce fluorine into HgBa2CaCu2O6þ (Hg1212) and HgBa2Ca2Cu3O8þ (Hg1223) compounds using BaF2 [94] or CuF2 [95] as a source of fluorine at 650–800 C, but no single phase samples were obtained. The superconducting phases, prepared in both cases, exhibit properties close to those for corresponding oxygenated materials. After treatment with XeF2 the reduced HgBa2CuO4.01 (Hg1201) phase (Tc ¼ 60 K) demonstrates an increase of Tc to its maximal value of 97 K, which is similar to the value achieved by optimal oxygen doping [96]. Increasing amount of fluorine above the optimal doping level results in a decrease of Tc and finally in a suppression of superconductivity due to overdoping. Both a and c unit cell parameters shrink concomitantly with the fluorine content indicating a contraction of Cu-O bonds due to increasing formal copper valence. Similar shrinkage of lattice parameters was observed also in the Hg1201, Hg1212 and Hg1223 samples fluorinated using CuF2, AgF2 or gaseous fluorine, where fluorinated samples demonstrated Tc as high as the oxygenated ones [97, 98]. Incorporation of fluorine into the superconducting Hg1201 phase was confirmed by the behaviour of the 19F NMR peak at the temperatures above and below Tc [99, 100]. Neutron powder structure refinement of the optimally doped (Tc ¼ 97 K) and overdoped (Tc ¼ 80 K) samples reveals that fluorine atoms are located in (1/2,1/2,0) position in the middle of Hg square mesh [96]. This was confirmed by the perturbed angular correlation technique, measuring the electric field gradients induced at 199mHg nuclei [101]. 19F NMR spectra revealed, however, that fluorine atoms occupy the (1/2,1/2,0) position only in an optimally doped sample, whereas at a higher doping level a part of the fluorine replaces oxygen at the apical positions of the CuO6 octahedra [102]. The occupancy factors for the (1/2,1/2,0) position in the fluorinated samples (n(F) ¼ 0.24(2) and n(F) ¼ 0.32(2)) are considerably higher than the occupancy factors n(O) for oxygenated samples with close Tc values (n(O) ¼ 0.124(9) and n(O) ¼ 0.19(1)). The parabolic dependencies of Tc versus extra anion concentration are similar for the fluorinated and oxygenated samples, but the curve for the fluorinated samples is shifted towards larger values (Figure 13.7). It indicates that the same doping level is achieved at a fluorine content level approximately twice higher than the oxygen

396

Functionalized Inorganic Fluorides 110

100

Tc, K

90

oxygen

fluorine

80

70

60

50 0.00

0.05

0.10 0.15 0.20 0.25 0.30 Extra oxygen / fluorine content

0.35

0.40

Figure 13.7 Tcvs. extra oxygen or fluorine content for Hg-1201 phase. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

content, i.e. oxygen, in agreement with its twice larger formal charge, creates twice as many holes as fluorine. This is also valid for the fluorinated Hg1223 phase with Hg0.8Ba2Ca2Cu3.2O8þ composition [103]. Due to partial replacement of Hg by Cu in the (HgO) layers, fluorine atoms occupy both the (1/2,1/2,0) position in the middle of Hg square mesh and the (1/2,0,0) position at the middle of the unit cell edge with the occupancies 27(1) % and 10(1) %, respectively. Similar occupancies of these positions (18 % and 8 %) were found in the oxygenated material with the same doping level [104]. Along with the carrier concentration, the in-plane Cu-Oeq bond length is an important parameter influencing the Tc value. The dependence of Tc versus a-parameter (or the doubled Cu-Oeq separation) for fluorinated and oxygenated Hg1201 samples can be fitted ˚ with the same parabolic-like curve with a maximum around 97 K at a ¼ 3.882 A (Figure 13.8). Thus the same Tc values are achieved in the fluorinated and oxygenated Hg1201 samples at the same doping level and with similar Cu-Oeq separations. The Cu-Oeq distance depends mostly on the copper oxidation state and practically does not vary with a change in the quantity of anions in the (Hg(O,F)) layers. In contrast, the apical Cu-Oap distance shows an almost linear decrease with increasing extra oxygen or fluorine concentration. The apical Cu-Oap distances differ significantly between the fluorinated and oxygenated samples with close Tc whereas the Cu-Oeq distance practically does not vary. This behaviour originates from a variation of the Hg-Oap distances. A shift of the Oap atoms away from Hg cations towards the Cu atoms is caused by elongating the Hg-Oap distance due to increasing the coordination number of Hg by twice larger amount of inserted fluorine. The shortening of the Cu-Oap distance results in anisotropic compression of the CuO6 octahedron, which is equivalent to an application of approximately 2 GPa of uniaxial

Properties of Cu- and Mn-based Oxyfluorides

397

110 oxygen fluorine

100

Tc, K

90

80

70

60

50 3.870

3.875

3.880

3.885

3.890

3.895

a, Å

Figure 13.8 The behavior of Tc vs. the a-parameter for the oxygenated and fluorinated Hg1201 phases. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

pressure along the c-axis. Application of 2GPa quasi-hydrostatic external pressure should enhance Tc by 4 K [105], but no difference in Tc was observed for the fluorinated and oxygenated optimally doped Hg-1201 samples. It seems that the enhancement of Tc under pressure is caused not by shortening the Cu-Oap distance but merely by the contraction of the in-plane Cu-Oeq distance. In the fluorinated Hg1223 phase the contraction of Cu-Oeq distance due to replacement of oxygen by larger amount of fluorine at the same doping level results in increasing transition temperature by 4 K in comparison with the optimally doped oxygenated sample [103]. The dependence of Tc versus the a parameter for three 140

Tc(K)

130

Hg-1223F

Hg-1223

Hg-1212

120 110 100

Hg-1201F Hg-1201

90 3.850 3.855 3.860 3.865 3.870 3.875 3.880 a-parameter (Å)

Figure 13.9 The dependence of optimal Tc vs. a-parameter for Hg-based superconducting Cu mixed oxides and oxyfluorides. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

398

Functionalized Inorganic Fluorides

members of the HgBa2Can1CunO2nþ2þ series and fluorinated Hg1201 and Hg1223 ˚ (Figure 13.9). One can conclude that phases is almost linear with dTc/da 1350 K/A the increase of Tc is caused by a compressive strain of the (CuO2) planes produced by chemical modification of the structure. It is interesting to note that a close ˚ value was observed due to the compressive epitaxial strain of the dTc/da 1.0 103 K/A La1.9Sr0.1CuO4 single-crystal thin film, which resulted in doubling Tc from 25 K to 49 K [106]. The value of dTc/da produced by such ‘chemical compression’ is much larger than that produced by application of external pressure. From crystallographic data of Armstrong et al. [107] and known dTc/dP 1.7 K/GPa for Hg-1223, the Tc vs the a parameter dependence ˚ . This value is caused by external pressure can be estimated as dTc/da ¼ 1.6 102 K/A much smaller than the dTc/da value caused by ‘‘chemical modification’’ (e.g. fluorination) of Hg-1223. It was concluded that chemical compression of the (CuO2) planes is a much more efficient factor for enhancing Tc than isotropic external pressure [97,103]. Most probably, noticeable buckling of the (CuO2) planes under high pressure and absence of such buckling in fluorinated samples are the origin of this effect. Application of 23 GPa pressure to fluorinated Hg1223 raised Tc to 166 K, the highest value ever observed [108].

13.3.3

Structural Rearrangements in Fluorinated Cuprates

The (R1,R2,A)2CuO4 (R1, R2 – rare-earth cations of different sizes, A – alkali-earth cation) cuprates are built of alternating (CuO2) planes and (R,A)2O2 blocks, which contain interstices suitable for the accommodation of overstoichiometric anions. These cuprates belong to three families known as T, T’ and T* phases. In the T-type structure (La2CuO4, La2xAxCuO4 (A – Sr, Ba), K2NiF4 structure type) the (R,A)2O2 blocks consist of the ((La,A)O)(O(La,A)) layers representing a motif of the NaCl structure. The Cu cations in the T phases are situated in oxygen octahedra. The cuprates of smaller rare-earth elements (R ¼ Pr – Gd) belong to the T’ family, in which the R2O2 blocks have fluorite-type structure with the (R)(O2)(R) layer sequence. In the T’ phases Cu cations adopt square planar coordination. Both NaCl- and fluorite-type blocks are present in the T* phases ((La,R,Sr)2CuO4, R ¼ Sm – Er)) according to the layer sequence: ðCuO2 Þ ððLa; SrÞOÞ ðOðLa; SrÞÞ ðCuO2 Þ ðR; LaÞ ðO2 Þ ðR; LaÞ ðCuO2 Þ The coordination polyhedron of the Cu cations in the T* phases is tetragonal pyramid. In all three types of structures the cations occupy the same positions, whereas the anion distribution in the (R,A)2O2 blocks is different (Figure 13.10). The R and A cations can be considered as composing a fragment of a close packed structure. In the ((R,A)O)(O(R,A)) block with the NaCl-type structure oxygens occupy octahedral interstices in the closepacked arrangement of the R and A cations, whereas in the fluorite-type (R)(O2)(R) blocks 50 % of the tetrahedral interstices are occupied. Thus, the anion positions in the (R,A)2O2 blocks, which are occupied in the T phases remain vacant in the T’ phases and vice versa. If extra anions enter into the structure, they can be accommodated in vacant interstices, which are the tetrahedral interstices in the NaCl-type blocks and octahedral interstices in the fluorite-type blocks. The separation between anions in the tetrahedral and octahedral ˚ , which should cause deformation of the structure containing interstices is only 2–2.15 A high amount of overstoichiometric anions in order to minimize the electrostatic repulsion. In the case of anion-deficient (R,A)2CuO4 phases anion vacancies can reside in the

Properties of Cu- and Mn-based Oxyfluorides

“NaCl”

399

-Cu -R,A -O -interstitial position

(CuO2)

2–2.15Å “CaF2”

Figure 13.10 Scheme of intergrowth between the (CuO2) planes, NaCl-type and fluorite-type blocks. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.) Sr2CuO3

Sr2CuO2F2+δ

CuO Sr2O2 CuO

-Sr -Cu -O -F

Sr2O2 CuO

Figure 13.11 The crystal structures of Sr2CuO3 and Sr2CuO2F2þ. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

(CuO2) planes, as it occurs in Sr2CuO3. It consists of alternating (SrO)(OSr) NaCl-like blocks and the (CuON) layers, so that the chains of corner-sharing CuO4 squares are present in this structure instead of (CuO2) layers [109] (Figure 13.11). In this case fluorination can cause significant rearrangements in the structure due to redistribution of oxygen and fluorine atoms. Upon fluorination of Sr2CuO3 the oxygen atoms from the (SrO)(OSr) blocks migrate to the (CuON) layers. As a result, in the Sr2CuO2F2þ oxyfluoride the Sr2O2 blocks are replaced by the Sr2F2þ blocks, the conducting (CuO2) planes are formed and Cu cations acquire octahedral coordination [26, 110] (Figure 13.11). This is also reflected by a ˚ ) along the direction perpendicular significant increase of the lattice parameter (by 0.9 A to the (CuO2) planes due to elongation of the apical distance in the octahedral environment of the Cu cations caused by Jahn-Teller effect. After fluorination, the Sr2F2þ blocks

400

Functionalized Inorganic Fluorides

contain extra fluorine atoms in the tetrahedral interstices ( 0.6), that results in holedoping level higher than necessary for superconductivity, but subsequent reductive annealing results in the optimal doping level ( 0.3) and a superconducting transition with Rc ¼ 46 K. Partial replacement of Sr by Ba in the Sr2xBaxCuO2F2þ increases Tc up to 64 K (x ¼ 0.6) [27]. Based on calculation of electrostatic lattice energy, the assumption was made that the configuration with the fluorine atoms at the apical positions in the octahedral copper environment is energetically more favourable in comparison with the placement of the fluorine atoms at the equatorial positions [26,110]. This conclusion was also supported by atomistic simulations [111]. Octahedral coordination of copper atoms with tentative CuO4F2 composition was detected by EXAFS, whereas XANES demonstrated that the doping is of p-type [49,112]. The structure of superconducting oxyfluoride Sr2CuO2F2þ belongs to the family of T-phases, consisting of alternating (CuO2) planes and the (SrF)(FSr) blocks with the NaCl-like structure. Being compared with the T phases, T’ phases require smaller A cations. It was demonstrated using the La2xNdxCuO4 solid solutions, where the transition from T to T’ structure occurs at x 0.5 due to difference in ˚ , CN ¼ 8) and Nd3þ (1.12 A ˚ , CN ¼ 8) [113]. In agreement with ionic radii of La3þ (1.18 A this tendency the Ca2CuO2F2þ oxyfluoride should adopt the T’-type structure [28]. Indeed, fluorination of Ca2CuO3 results in fluorine insertion into tetrahedral interstices of the Ca2F2þ block, and overstoichiometric anions are located in the octahedral interstices, completing coordination environment of a part of the Cu atoms to tetragonal pyramid [114]. In contrast to its Sr-based analogue, Ca2CuO2F2þ does not demonstrate superconducting properties. A presence of the Cu atoms in two different coordinations (square planar and tetragonal pyramidal) due to partial occupation of the octahedral interstices in the Ca2F2þ blocks is, perhaps, at the root of the absence of superconductivity because it results in local buckling of the (CuO2) planes and charge carrier localization. Anion rearrangement, similar to that observed in fluorinated Sr2CuO3, was also induced by fluorination in Pb2Sr2YCu3O8 [37]. The layer sequence in this compound is (CuN2) (PbO) (OSr) (CuO2) (YN) (CuO2) (OSr) (PbO) (CuN2) . In the (CuN2) layer the Cu cations are situated in a dumbbell coordination being surrounded by two apical oxygens and oxygen vacancies in all equatorial positions. Treatment of Pb2Sr2YCu3O8 with NH4F results in exsolution of PbF2 and a formation of the PbSr2YCu3O6F4 oxyfluoride. As in Sr2CuO2F2þ, apical oxygen anions forming dumbbells around the Cu cations migrate to equatorial positions, being replaced by fluorine anions, which leads to CuO4F2 octahedral coordination around Cu. Due to transformation of the dumbbells into Jahn-Teller elongated octahedra, the c-parameter of the fluorinated ˚ ) is much larger than c-parameter of the initial PbSr2YCu3O6F4 phase (17.062 A ˚ ). In the fluorinated compound the Pb positions are only Pb2Sr2YCu3O8 phase (15.733 A 50 % occupied due to the release of PbF2. A formation of this admixture with high lattice energy was considered as a driving force for the fluorination reaction. In contrast to A2CuO3 (A ¼ Ca, Sr, Ba), La2CuO4 with the T-type structure does not require anion redistribution to demonstrate superconducting properties because (CuO2) planes already exist in this structure (Figure 13.12). The required hole concentration can be created either by heterovalent replacement in the cation sublattice (La2xAxCuO4 (A ¼ Sr,Ba) solid solutions with Tc ¼ 35 – 40 K [115, 116]) or by insertion of extra oxygen anions through treatment under elevated oxygen pressure [117, 118], electrochemical oxidation [119] or treatment with KMnO4 solution [120]. Fluorine insertion into the La2CuO4 structure also causes the appearance of superconductivity [29, 35, 40, 121,

Properties of Cu- and Mn-based Oxyfluorides

401

CuO2 La2O2

Figure 13.12 The crystal structure of La2CuO4. The Cu cations are situated in octahedra, the La cations are shown as spheres

122]. At the first stage of fluorination with pure or diluted F2 or with XeF2 at a relatively low temperature of 150–200 C, the superconducting samples with Tc 35 – 40 K are obtained. Insertion of small amount of fluorine, up to La2Cu(O,F)4.18, composition is accompanied by an increase of orthorhombic distortion and slight increase in the c-parameter [29]. Overstoichiometric fluorine anions are located in the tetrahedral interstices in the NaCl-type La2O2 blocks [121]. Structural consequences of fluorine and oxygen insertion are nearly the same [123]. The degree of orthorhombic distortion depends on the amount of inserted anions being very close for the samples with similar ˚ , b ¼ 5.427 A ˚ , c ¼ 13.194 A ˚ for La2CuO4F0.18 [121], anion content (a ¼ 5.328 A ˚ , b ¼ 5.423 A ˚ , c ¼ 13.222 A ˚ for La2CuO4.19 [124]). Fluorination of the a ¼ 5.332 A La2xAxCuO4 (A ¼ Sr, Ba) solid solutions affect the superconducting properties only for compounds with small x values, also reproducing the effect of oxygen insertion. Superconductivity appears after fluorination of the La2xSrxCuO4 (x < 0.1, Tc ¼ 36 – 40 K) [121, 125] and La2xBaxCuO4 (x < 0.15, Tc ¼ 32 42 K) solid solutions [122, 126], whereas Tc for the initial and fluorinated samples with larger x values remains unchanged. Increasing temperature of fluorination with F2 or XeF2 up to 230–250 C results in a nonsuperconducting oxyfluoride with tetragonal unit cell and lattice parameters reflecting significant expansion of the structure in the ab plane and contraction along the ˚ and c axis [127–129]. The K2NiF4 subcell parameters in the (CuO2) plane are as 3.802 A ˚ for the initial and fluorinated samples, respectively. Electron diffraction as 4.038 A reveals a superstructure with doubling of both a and c parameters of the K2NiF4 subcell, which is caused by alternation of vacant and filled tetrahedral interstices. Such ordering can be accompanied by cooperative tilting of the CuO6 octahedra in order to maximize the distances between the apical oxygen anions and anions in the tetrahedral interstices. Fluorination of La2CuO4 with XeF2 at higher temperatures (300–400 C) occurs mainly as anion exchange and leads to the La2CuO3.6F0.8 oxyfluoride with the copper oxidation state þ2 [129]. The basic vectors of the monoclinic structure of this compound are related to the basic vectors of the K2NiF4 subcell, such as ak ¼ 2(aK2NiF4 þ bK2NiF4) þ cK2NiF4, bk ¼ aK2NiF4 þ bK2NiF4, ck ¼ ½(3aK2NiF4 þ 3bK2NiF4 - cK2NiF4). In the structure model, deduced from transmission electron microscopy data, fluorine atoms occupy 40% of tetrahedral interstices in an ordered manner, removing part of the oxygen atoms from

402

Functionalized Inorganic Fluorides

cm bm

am F

F

F

F

Figure 13.13 The crystal structure of La2CuO3.6F0.8 and scheme of cooperative rotation of the Cu coordination polyhedra. The Cu cations are situated in octahedra, the La cations are shown as large spheres, the F anions are shown as small spheres.

the (LaO) layers and transforming part of the CuO6 octahedra into CuO5 square pyramids (Figure 13.13). The remaining octahedra are cooperatively rotated maximizing the distance between the apical oxygen atoms and interstitial fluorines. Thus, the blocks with filled tetrahedral interstices and empty octahedral interstices (fluorite fragments) are present in the structure along with the NaCl-like fragments of the initial structure, but in contrast to the T* structure, these blocks alternate not only along [001]K2NiF4, but also along [110]K2NiF4. La2CuO4 can be also considered as the n ¼ 1 member of the Ruddlesden-Popper homologous series with general (R,A)nþ1CunO2nþ2þ (R ¼ rare-earth cation, A ¼ Ca, Sr) composition. In the structure of the n ¼ 2 member R2xA1þxCu2O6þ the (CuO2) planes are separated by the (AN) layer (Figure 13.14): ðCuO2 Þ ðOðR;AÞÞ ððR;AÞOÞ ðCuO2 Þ ðANÞ ðCuO2 Þ ððR;AÞOÞ ðOðR;AÞÞ ðCuO2 Þ-: The cation position in the (AN) layers can be jointly occupied by rare-earth and alkaliearth cations if A ¼ Sr. The Cu cations have tetragonal pyramidal oxygen coordination. In contrast to the n ¼ 1 member, in the structure of the second member there are two possible positions for extra anions: in the tetrahedral interstices between the (RO) and (OR) layers and in the (AN) layer, that will complete the Cu coordination to octahedron. The behaviour

Properties of Cu- and Mn-based Oxyfluorides

403

A CuO2 (R,A)2O2 CuO2 A n=2

n=3

Figure 13.14 The crystal structures of the n ¼ 2 and n ¼ 3 (R,A)nþ1CunO2nþ2þ RuddlesdenPopper type compounds

of the n ¼ 2 members under fluorination differs from behaviour of La2CuO4. Treatment of La1.9Sr1.1Cu2O6.05 and La1.9Ca1.1Cu2O5.95 was done with F2, NH4F and CuF2 [31]. The reaction with F2 proceeds via fluorine insertion, while reaction with NH4F involves fluorine substitution for oxygen. The reaction with CuF2 at low temperature probably occurs as both insertion and substitution, while at higher temperatures (>300 C) the substitution process dominates. The fluorination results first in increasing a and c parameters of the tetragonal unit cell. At this stage fluorine is supposed to enter into tetrahedral interstices of the NaCl-like blocks. A larger degree of fluorination introduces orthorhombic distortion into the Sr-containing structure with decreasing lattice parameters ˚ to 21.68 A ˚ . Such an in the ab plane and drastic increase of the c parameter from 20.048 A increase is related to the occupation of vacant anion positions in the (AN) layers by fluorine atoms and formation of octahedral coordination for at least a part of the Cu atoms with elongated Cu – (O,F) apical bonds due to Jahn-Teller effect. Such effect was not observed in fluorinated La1.9Ca1.1Cu2O5.95, probably because of a smaller distance between the (CuO2) planes due to smaller size of the Ca cations, which creates sterical difficulties for fluorine insertion into the (AN) layers. Both fluorinated La1.9Sr1.1Cu2O6.05 and La1.9Ca1.1Cu2O5.95 are not superconducting. Superconducting n ¼ 2–5 members of the Ruddlesden-Popper series were prepared using the high-pressure, high-temperature technique. The Sr2CaCu2O4.6F2 and Sr2Ca2Cu3O6.2F3.2 samples [15,16] demonstrate bulk superconductivity with Tc ¼ 99 K and 111 K, respectively, although they contain significant amount of admixture phases. For higher members of this homologous series Sr2Ca3Cu4O8þF2 and Sr2Ca4Cu5O10þF2 Tc was lower (40–95 K) [130]. For the Ba-based series Ba2Can1CunO2n(O,F)2 the highest Tc for n ¼ 2, 3, 4 and 5 were 90, 120, 105 and 90 K, respectively [17,18]. The Sr2RCu2O5F (R ¼ Y, La, Nd, Sm, Gd, Dy, Er and Yb) compounds are obtained under pressure and are nonsuperconducting because of VCu ¼ þ2 [14]. Heterovalent replacement of Nd by

404

Functionalized Inorganic Fluorides

Ca results in the Sr2Nd0.2Ca0.8Cu2O5F compound with Tc ¼ 85 K. It is assumed that in these phases the fluorine anions are located in the NaCl-type blocks, forming the (SrF)(FSr) or (SrO0.5F0.5)(F0.5O0.5Sr) layers. The effect of fluorination on the structure and properties of the T’-phases was most comprehensively studied with fluorine-containing derivatives of Nd2CuO4 but the majority of studies deal with the Nd2CuO4xFx, where n-type superconductivity was found with Tc ¼ 27 K for x ¼ 0.3 [9]. As was discussed in Section 13.3.1, a replacement of oxygen with an equivalent amount of fluorine in the fluorite-type Nd2O2 slabs does not significantly alter the crystal structure. Insertion of fluorine into the Nd2CuO4 structure results in a new compound with the idealized Nd2CuO3F2 composition [131]. This compound was found using transmission electron microscopy in the samples of Nd2CuO4, fluorinated with XeF2 at ˚ , b 5.5 A ˚ , c 5.8 A ˚, 300–400 C. In the monoclinic structure of Nd2CuO3F2 (a 12.96 A 92) the anions in the tetrahedral interstices are arranged into chains running along the c-axis and alternating along the b-axis with the chains of vacant interstices (Figure 13.15). Some of the anions occupy the apical position in the coordination environment of the Cu cations. A formation of the octahedral coordination environment for the Cu cations results in ˚ in the initial material to 12.96 A ˚ in the an increase of the c-parameter from 12.167 A fluorinated phase. The Cu(O,F)6 octahedra are cooperatively tilted to maximize the distances between apical anions and anions in the tetrahedral interstices. The copper oxidation state in the fluorinated samples was not high enough to achieve p-type superconductivity.

a c b

Nd

O, F

Figure 13.15 The crystal structure of Nd2CuO3F2. The Cu cations are situated in octahedra. The vacant tetrahedral interstices are shown as empty squares

A first stage of fluorination with XeF2 of Nd2CuO3.5, an anion-deficient derivative of Nd2CuO4, results in a superconductor with Tc ¼ 6 – 11 K [132]. Nd2CuO3.5 consists of fluorite-type Nd2O2 slabs and (CuO1.5N0.5) layers with Cu atoms in dumbbell and strongly distorted square-planar coordinations [133], so that there are no conducting (CuO2) planes in this structure (Figure 13.16). Fluorination at 200 C leads to an increase of the formal

Properties of Cu- and Mn-based Oxyfluorides

405

copper valence and to a restoration of the T’-type structure. Fluorine replaces oxygen in the Nd2O2 slab and the released oxygen atoms migrate into the (CuO1.5N0.5) layers thus forming the (CuO2) planes (Figure 13.16). This rearrangement shows a close resemblance to that previously found in fluorinated A2CuO3 (A ¼ Ca, Sr). Increasing fluorination temperature results in the Nd2Cu(O,F)5 with the monoclinic structure of Nd2CuO3F2, as obtained after fluorination of Nd2CuO4 (Figure 13.16).

F

(CuO2) F

13.0Å

11.9Å 12.5Å

Nd2CuO3.5

Nd2CuO3.5F -Nd

-Cu

-O

Nd2Cu(O,F)5 -F, O

Figure 13.16 Scheme of the structure transformations under fluorination of Nd2CuO3.5. Structure of the reduced compound Nd2CuO3.5 with a linear coordination for copper; the fluorinated superconducting Nd2CuO3.5Fx phase with the (CuO2) planes and T’ structure, fluorine occupies the tetrahedral positions in the fluorite type (Nd2(O,F)2) block; highly fluorinated Nd2Cu(O,F)5 phase with oxygen and fluorine distributed in an ordered manner over both the octahedral and tetrahedral interstices, forming chains. (Reprinted with permission from [132] Copyright (2003) American Chemical Society.)

Nd2CuO3F2 is to some degree similar to the La2CuO3.6F0.8 structure: both can be considered as intermediate between the T and T’ structures and resemble the structural organization of the T* phases, where the NaCl- and fluorite-type blocks alternate along the c-axis. However, in both Nd2CuO3F2 and La2CuO3.6F0.8 occupied and vacant tetrahedral and octahedral interstices alternate not only along the c-axis of the tetragonal subcell but also in the ab-plane. In both cases the ratio between the parameters of the tetragonal subcells c/3a 1.08 falls into the stability region of the T* phases (c/3a 1.15 for T-phases, c/3a 1.06–1.08 for T* and c/3a 1.02 for the T’ phases [134]). The different effects of fluorination on the (R,A)2O2 blocks with the NaCl- and fluoritetype structures can be illustrated by the behaviour of the T* phases. The results of fluorination of the (R1, R2, M)2CuO4 do not strongly depend on the cation composition of the initial materials. In the La1.25Dy0.75CuO3.75F0.5 [45] and LaBa0.2Gd0.8CuO3.55F0.8 [135] crystal structures, obtained by fluorination with ZnF2 and CuF2, excessive anions are located in the tetrahedral interstices of the NaCl-type block, whereas the fluorite block remains unchanged. Anion – anion repulsion and/or decreasing copper oxidation state ˚ to 1.942–1.962 A ˚ and slight result in increasing interplanar Cu-Oeq distances from 1.933 A increase of the c-parameter. The LaBa0.2Gd0.8CuO3.55F0.8 retains the tetragonal symmetry p p of the initial material, whereas La1.25Dy0.75CuO3.75F0.5 is orthorhombic with the 2 2 superstructure in the ab-plane. In both cases the apical anions of the Cu(O,F)5 pyramids

406

Functionalized Inorganic Fluorides

are displaced from their ideal positions, most probably due to their interaction with interstitial anions [136]. In the LaHo0.75Sr0.25CuO3.9 R*-phase, fluorinated with XeF2, inhomogeneous fluorine distribution results in two co-existing phases [137]. In the orthorhombic phase with relatively low fluorine content the structure of the fluorite-type blocks is not affected, whereas the NaCl-type blocks acquire a structure similar to that in with a larger fluorine content has a monoclinic unit cell with La2CuO3.6F0.8. The phase ˚ atp 2, cm 12.84 A ˚ ct, 91.4 (t belongs to the tetragonal subcell). am bm 5.5 A In this phase, as in the case of the T’-phase, larger cm value in comparison with ˚ reflects that extra anions partially occupy the apical positions of copper ct ¼ 12.553 A coordination environment. The anions are equally distributed among the tetrahedral and octahedral interstices of the initially different (La0.75Sr0.25O)(OLa0.75Sr0.25) (NaCl-type) and (La0.25Ho0.75)O2(La0.25Ho0.75) (fluorite-type) blocks. Thus, at higher fluorination degree, the structural changes occur in both types of blocks. Vacant and occupied tetrahedral sites in these blocks alternate in a ‘chess-board’ manner. The monoclinic distortion arises from cooperative displacement of the R- and A-cations due to interaction with ordered anions in the tetrahedral interstices. From the results of structural investigations on the T, T’ and T* phases several common trends can be derived [138]. Insertion of fluorine atoms does not directly affect the (CuO2) planes, but it can cause distortions due to cooperative tilting of the Cu coordination polyhedra. Fluorine atoms enter into the (R,A)2O2 blocks. Fluorine enters more easily into the NaCl-type blocks, whereas fluorination causes substantial structural changes in the fluorite-type blocks. At a high degree of fluorination the anion rearrangement occurs in both blocks resulting in the structures where short fragments of the NaCl and fluorite structures alternate along the direction perpendicular to the (CuO2) planes, as well as in the plane parallel to the (CuO2) planes. Such structures can be considered as intermediate between the structures of the T and T’ phases and are to some degree similar to the T* structures. Completing copper coordination environment to an octahedron is at the origin of the significant increase of the c-parameter in the fluorinated T’ and T* phases. In the (R,A)2O2 blocks, where R and A are the rare-earth and alkali-earth cations, respectively, the bonding between the cations and neighbouring anions is mostly ionic. Thus, in the NaCl-type (R,A)2O2 blocks the bonds within the ((R,A)O) layers are as strong as the bonds linking the neighbouring ((R,A)O) layers. The NaCl-like Bi2O2 blocks are also present in the structures of Bi-based layered cuprates Bi2Sr2Can1CunO2nþ4þ. Structural modulations are intrinsic in the (BiO) layers, associated with a mismatch between the (BiO) layers and perovskite-like fragment. This mismatch is reduced via strong shifts of atoms from their ideal sites and insertion of additional oxygen anions in the bismuth-containing layers [139]. In contrast to ionic type of interactions in the (R,A)2O2 blocks, only weak Van der Waals interaction occurs between two adjacent (BiO) layers. Due to a weak link between the (BiO) layers, different halogen-containing species can be intercalated into the Bi-based layered cuprates (I2, Br2, HgBr2, LiI3) [140–144]. The iodine atoms in the n ¼ 2 Bi2212 and n ¼ 3 Bi2223 structures form an intermediate layer between the (BiO) and (OBi) layers, resulting in significantly increasing separation between these layers. The structural modulations of the (BiO) layers remain because weak interaction with the iodine atoms does not significantly affect the atomic arrangement in these layers. On the other hand, the Bi-F bond is more ionic in comparison with other halogens and fluorine insertion can result in substantial changes in the structure of the (BiO) layers. In

Properties of Cu- and Mn-based Oxyfluorides

407

the superconducting oxyfluoride Bi2Sr2Ca2Cu3O8F4 (Tc ¼ 75 K) all oxygen atoms in the (BiO) and (OBi) layers are supposed to be replaced by fluorine and extra fluorine anions are located in the tetrahedral interstices between the layers. [43,145]. In this structure in the –(BiF)-(F2)-(FBi)- blocks the Bi atoms have coordination environment similar to that ˚ in in BiF3. The separation between the (BiF) and (SrO) layers is increased by 0.9 A comparison with the separation between the (BiO) and (SrO) layers in the initial Bi2223 phase. At the same time, the distance between the (BiF) and (FBi) layers is decreased by ˚ , probably due to bonding between the Bi atoms and interstitial fluorines. 0.7 A Fluorination of the Bi2201 phase Bi2Sr1.6La0.4CuO6.33 with XeF2 results in an increasing c-parameter of the unit cell and suppression of structural modulations [146]. This reflects the fact that the inserted fluorine changes the structure of the (BiO) layers. The inserted anions are located in the tetrahedral interstices of the Bi2O2 block, so that it is transformed into the Bi2O2xF2þx block, where the Bi atoms are located in a capped square antiprism with four short Bi-F bonds and five longer Bi-O bonds (Figure 13.17). Blocks with a similar structure were found earlier in the members of Aurivillius homologous series, such as Bi2NbO5F [147]. A comparison of interlayer distances in the Bi2O2 blocks of the initial Bi2201 phase and in the Bi2O2F2 block of the Aurivillius phase explains the ˚ in the initial Bi2201 phase to 26.18 A ˚ in the increase in the c-parameter from 24.35 A

Bi (La,Sr) (O,F)

Figure 13.17 The crystal structure of the fluorinated Bi2201 phase. The Cu cations are situated in octahedra. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

408

Functionalized Inorganic Fluorides - Bi -O -F

Bi2201

4.5Å

2.6Å

3.2Å

3.7Å

- Nb or Cu - Sr

Bi2NbO5F

Figure 13.18 Comparison of the crystal structures of Bi2201 phase and the Aurivillius-type Bi2NbO5F phase. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

fluorinated Bi2201 phase (Figure 13.18). A distance between two layers of the Bi cations ˚ in the initial above and below the tetrahedral interstices decreases from 3.18–3.20 A ˚ Bi2201 phase to 2.49–2.56 A in the fluorinated Bi2201 phase due to electrostatic interaction between the fluorine anions and the Bi3þ cations. On the other hand, the oxygen anions in the (BiO) layers are repelled by the interstitial anions. Thus, in the initial Bi2201 structure, the (BiO) layers are nearly flat and the projected distance along the c-axis ˚ . In the Aurivillius phase this between the Bi and O atoms does not exceed 0.24–0.26 A ˚ distance increases to 0.92–1.07 A, resulting in ‘splitting’ of the layer into two layers consisting of the Bi and O atoms, respectively. In spite of decreasing distance between ˚, the layers of the Bi cations, overall expansion of the Bi2O2F2 blocks is about 0.7–1.0 A ˚ resulting in the increase of the c-parameter of 1.4–2.0 A. Increasing coordination number of the Bi cations causes elongation of the Bi-O distances and expansion of the Bi2O2F2 block in the ab-plane, eliminating the mismatch between the Bi2O2F2 blocks and (CuO2) planes, which suppresses structural modulations. 13.3.4

Fluorination of Nonsuperconducting Cuprates

The discovery of the possibility of rearranging oxygen and fluorine atoms with the formation of the (CuO2) planes in fluorinated Sr2CuO3 and Nd2CuO3.5 opened new possibilities to search for new superconductors among anion-deficient cuprates. These cuprates do not contain the (CuO2) planes, but the planes could be formed due to anion rearrangement caused by fluorination. One of the simplest copper oxyfluorides with possible superconducting properties would be the perovskite-like (La,Sr)CuO2F solid

Properties of Cu- and Mn-based Oxyfluorides

409

solutions, where the required copper oxidation state could be achieved by changing the La/Sr ratio. Such compound could be synthesized by soft fluorination of the aniondeficient perovskites with (La,Sr)/Cu ¼ 1. There is a large variety of perovskite-type cuprates with three-dimensional frameworks of corner-sharing copper-oxygen polyhedra, which differ mainly by ordering patterns of oxygen atoms and vacancies: the columns of oxygen atoms running along the [001] direction of the perovskite subcell are removed in an ordered manner resulting in decreasing coordination number of the copper atoms from 6 to 5, 4 or 2 [148]. The (Lap1xSrx)8Cu8O20 solid solutions with tetragonal structure with lattice parameters a ap2 2, c ap belong to this structural family [149, 150] (Figure 13.19). Fluorination of La6.5Sr1.5Cu8O19.65 with XeF2 produces a tetragonal perovskite phase with the La0.813Sr0.187Cu(O,F)3 composition and lattice parameters a < c ap [151]. In the La0.813Sr0.187Cu(O,F)3 structure the equatorial anion positions in the copper coordination environment are fully occupied and anion vacancies are located in the apical positions. Octahedral coordination is formed for at least part of the Cu cations with elongated apical ˚ and shorter equatorial ones of 1.896 A ˚ , that is also reflected in a distances of 2.026 A tetragonal distortion of the perovskite subcell with c > a. The La0.813Sr0.187Cu(O,F)3 oxyfluoride is not a superconductor.

La6.5Sr1.5Cu8O20-δ

La,Sr

La0.813Sr0.187Cu(O,F)3-δ

1.90Å

Cu

La,Sr

2.03Å

Figure 13.19 The crystal structures of La6.5Sr1.5Cu8O20 and La0.813Sr0.187Cu(O,F)3. (Reprinted with permission from Russ. Chem. Rev., A. M. Abakumov et. al., 71, 5, 383–400 Copyright (2004) Royal Society of Chemistry.)

410

Functionalized Inorganic Fluorides

Perovskite-like cuprates with the brownmillerite-type structure RACuGaO5 (R ¼ La, Pr, Nd; A ¼ Ca, Sr) (Ga1201) [152, 153] provide a structure suitable p for superconductivity. The orthorhombic brownmillerite structure (a 4ap, b c ap 2) is represented as a sequence of layers (Figure 13.20): ðCuO2 Þ ðR0:5 A0:5 OÞ ðGaONÞ ðR0:5 A0:5 OÞ ðCuO2 Þ The Ga and Cu cations possess the tetrahedral and octahedral oxygen coordination, respectively. Within the (GaON) plane the GaO4 tetrahedra are shared by common vertexes into chains running along the c-axis. Ordering of oxygen anions and vacancies in the (GaON) layers distorts the atomic arrangement in the (CuO2) planes leading to two ˚ ) and two long (1.96 A ˚ ) Cu – Oeq distances [154,155]. Significant difference short (1.89 A in the Cu – Oeq bond lengths is the most possible reason why these compounds do not become superconductors even at appropriate doping level. Indeed, the La1xSr1þxCuGaO5 (0 £ x £ 1.2, max. VCu ¼ þ2.2) solid solutions do not show superconducting transition at any x value [156]. However, in the structure of the next member of brownmillerite-type homologous series Ga1212 Y0.6Ca0.4Sr2Cu2GaO7 the Cu cations are situated in square ˚ h 1.93 A ˚ ), and this compound is pyramids with almost identical Cu – Oeq distances (1.92 A a superconductor with Tc ¼ 44 K [157]. In order to induce superconductivity in Ga1201 it is necessary not only to achieve the optimal doping level, but also to suppress the distortions caused by anion vacancy ordering in the (GaON) layers, that was expected to be done by fluorine insertion into vacant positions in these layers. Treatment of LaACuGaO5 (A ¼ Ca, Sr) with XeF2 decreases orthorhombic distortion up to a transformation to a fluorinated phase with a tetragonal symmetry (for A ¼ Sr) with a ap, c 2ap [158]. The copper oxidation state in the fluorinated samples depends strongly on the conditions of the fluorination reaction. Fluorination with a large excess

LaSrCuGaO5

LaSrCuGaO5-xF2x-δ

- La(Sr)

CuO2 La0.5Sr0.5O GaO La0.5Sr0.5O CuO2

Figure 13.20

CuO2 Ga(O,F)δ CuO2

The crystal structures of LaSrCuGaO5 and LaSrCuGaO5xF2x

Properties of Cu- and Mn-based Oxyfluorides

411

of XeF2 at a relatively low temperature (300 C) for a long time (150 h), occurs via combined anion exchange and fluorine insertion resulting in VCu ¼ þ2.16: • LaSrCuGaO5 þ ðx þ y=2ÞXeF2 !LaSrCuGaO5 x F2x þ y þ x½O þ ðx þ y=2ÞXe . Increasing temperature to 400 C leads to slight reduction of VCu because the anion exchange dominates at higher temperature: LaSrCuGaO5 þ ðx =2ÞXeF2 ! LaSrCuGaO5 x F2x þ x½O þ ðx =2ÞXe; VCu ¼þ1:95 In the fluorinated phase fluorine atoms together with the oxygen atoms occupy positions in the (GaON) layers, increasing coordination number of the Ga cations to five and six. In the tetragonal structure of the fluorinated phase the (CuO2) planes are flat and all the Cu – Oeq distances are equivalent (Figure 13.20). However, remaining local distortions do not allow the fluorinated brownmillerite to be a superconductor.

13.4

Fluorination of Manganites

The interest in manganites is related to great extend to colossal magnetoresistance (CMR), commonly associated with three-dimensional perovskite-like oxides R1xAxMnO3– (R is a rare-earth cation, A is a mono- or divalent metal) [159–164]. Magnetic ordering and transport properties of R1xAxMnO3– are sensitive to the structural features influenced by the average (R,A) -cation radius and the Mn oxidation state. Electronic concentration, which can be varied by anion doping or heterovalent cation replacements, is relevant for realization of exchange interactions between the Mn cations, which influences the conducting and magnetic behaviour. The magnetic phase diagrams and charge-ordering phenomena in complex manganites are determined by a conductive band width, which depends in part on Mn-O-Mn in-plane bond angle and can be varied by changing r(R,A)/rMn ratio [165, 166]. According to the double exchange mechanism, the strong Hund’s coupling of eg and t2g electrons results in FM alignment of the t2g electrons through delocalized eg ones (Mnþ3 (t2g3eg1) ! Mnþ4 (t2g3eg0)), i.e. ferromagnetism in manganites is associated with metallic behaviour. The r(R,A)/rMn ratio determines a degree of distortion of the perovskite framework in order to allow small R and A cations to adopt the coordination number less than 12. This distortion is realized in cooperative rotations and tilts of the MnO6 octahedra, resulting in symmetry decrease of the initially cubic perovskite aristotype. The buckling of the octahedral framework prevents delocalization of the eg electrons and stabilizes the charge ordered state. The perovskites with large r(R,A)/rMn generally show FM metallic behaviour, whereas the insulating charge-ordered states are usually observed for the compounds with smaller rA/rMn ratios. It is of interest to attempt tuning the magnetic and conducting properties of manganites using fluorination. In this case the heterovalent replacement in the anion sublattice may allow to explore the magnetic properties as a function of the doping level at fixed size of the (R,A) cations. p p Sr2Mn2O5 is an anion-deficient perovskite with an orthorhombic a ap 2, b 2ap 2, c ap superstructure due to ordering of anion vacancies [167] (Figure 13.21). One quarter of the [001] rows of the oxygen atoms is missed, resulting in a network of corner-sharing distorted tetragonal pyramids MnO5. This compound is antiferromagnetic (AFM)

412

Functionalized Inorganic Fluorides

Sr

Figure 13.21 The crystal structure of Sr2Mn2O5. The Mn cations are located in the distorted tetragonal pyramids

with TN ¼ 380 K. Fluorination of Sr2Mn2O5 with XeF2 produces a mixture of SrF2 and ˚ , c ¼ 3.989(2) A ˚ ) with the Sr2Mn2O5F tetragonally distorted perovskite (a ¼ 3.8069(8) A structure, similar to that of La0.813Sr0.187Cu(O,F)3 (see Figure 13.19, bottom) [60]. The tetragonal distortion of the perovskite structure obviously originates from the Jahn–Teller distortion of the Mn(O,F)6 octahedra. The octahedra are ‘apically elongated’ with four ˚ and two longer apical Mn–Oap ones short equatorial Mn–Oeq distances of 1.903 A ˚ ). The sample shows a ferromagnetic upturn in susceptibility at 55 K in agreement (1.995 A with the double exchange interaction between d4 Mn3þ and d3 Mn4þ giving rise to ferromagnetic coupling. In spite of large diversity of structures due to a complex interplay between geometrical and electronic effects, the perovskite-like R1xAxMnO3– oxides retain a three-dimensional net of the Mn-O bonds. The Ruddlesden-Popper series (R,A)nþ1MnnO3nþ1 provide an intermediate case of dimensionality from D ¼ 3 in perovskites (n ¼ 1) down to D ¼ 2 in the (R,A)2MnO4. These compounds are regarded as potential CMR materials working at fields much lower than 3D perovskites [168,169]. Fluorine-containing derivative of the n ¼ 2 Ruddlesden-Popper type manganite La1.2Sr1.8Mn2O7 [41] was obtained by interaction with CuF2, which results in incorporation of 2F per formula unit. Fluorine anions are located in tetrahedral interstices of the NaCl-like block. In comparison with the initial material, La1.2Sr1.8Mn2O7F2 has slightly smaller a-parameter of the tetragonal unit cell and significantly larger c-parameter. Contraction of the equatorial Mn-Oeq distances from ˚ to 1.885 A ˚ is in agreement with the increasing Mn oxidation state. Thickness of the 1.96 A ˚ in La1.2Sr1.8Mn2O7F2. ˚ in La1.2Sr1.8Mn2O7 to 7.68 A NaCl-block increases from 6.18 A Similar structural changes were also observed after fluorination of R1.2Sr1.8Mn2O7 (R ¼ Pr, Nd, Sm, Eu, and Gd) [59]. Enlargement of the c-parameter was also observed in fluorinated Ruddlesden-Popper type cuprates, but the structural reasons for this enlargement seem to be different. In cuprates vacant anion positions in the (AN) layers are occupied by fluorine resulting in apically elongated Jahn-Teller distorted Cu(O,F)6

Properties of Cu- and Mn-based Oxyfluorides

413

octahedra. In La1.2Sr1.8Mn2O7F2 fluorination does not directly affect the coordination environment of the Mn cations. Moreover, the Jahn-Teller deformation is not possible due to VMn > þ4. The increase in c-parameter in La1.2Sr1.8Mn2O7F2 is caused by expansion of the NaCl-block due to repulsion between interstitial anions and apical anions of the MnO6 octahedra. A staged fluorine insertion into LaSrMnO4 [67,68] and La1.2Sr1.8Mn2O7 [64] was achieved by a solid-state reaction of the fluorinated manganite and its oxide prototype. LaSrMnO4F1.7 reacts with LaSrMnO4 at 300 C producing LaSrMnO4F. La1.2Sr1.8Mn2O7F was obtained in the same way. In both LaSrMnO4F and La1.2Sr1.8Mn2O7F compounds only 50 % of the NaCl-type blocks contain fluorine in the tetrahedral interstices, and the blocks with and without fluorine alternate along the c-axis (Figure 13.22). Similar to the La1.2Sr1.8Mn2O7F2 structure, the fluorine-containing NaCl-blocks are wider than the blocks without fluorine, which also causes a noticeable increase in the c-parameter. LaSrMnO4F and La1.2Sr1.8Mn2O7F do not demonstrate longrange magnetic ordering because for Mn4þ with t2g3eg0 only weak p-type t2g-O(2p)-t2g interactions are possible.

(La,Sr)2O2 MnO2 (La,Sr)2O2F2

Figure 13.22 The crystal structure of LaSrMnO4F. The Mn cations are situated in octahedra, the La and Sr cations are shown as grey spheres, the F anions – as dark grey spheres

An alternative way towards layered Mn-based perovskites comprises a replacement of Mn by another B-cation with the crystal chemistry properties different from those for the Mn-cations that favours the ordered layered arrangement. The brownmillerite-type Sr2MnGaO5þ manganite is formed due to incorporation of layers of tetrahedrally coordinated cations into the perovskite framework [170,171] (see the brownmillerite structure in Figure 13.20). The oxygen content in Sr2MnGaO5þ brownmillerite can be varied in the -0.03 < < 0.13 and 0.41 < < 0.505 ranges [172,62]. The reduced ( 0) Sr2MnGaO5þ compound is AFM ordered with G-type magnetic structure where the Mn magnetic moments are aligned perpendicular to the (MnO2) planes [173–175]. In the oxidized ( 0.5) Sr2MnGaO5þ the extra oxygen atoms in the (GaO1þ) layers alter the magnetic structure to be AFM of C-type [173, 175, 176]. The transition from G to C-type magnetic structure occurs due to a strong diagonal 180 AFM interaction between the Mn4þ-cations in adjacent layers through the additional oxygen atoms in (GaO1þ) layers. The anion content in the Mn-based brownmillerites can be smoothly changed also by fluorination with XeF2

414

Functionalized Inorganic Fluorides

[61,62]. Starting from oxidized Sr2MnGaO5.5 compound, the fluorination reaction occurs at 500–600 C mainly as anion exchange, resulting in a reduction of the Mn valence: Sr2 MnGaO5:5 þ ð1 þ xÞ=2 XeF2 ! Sr2 MnGaO5 x F1 þ x þ ð0:5 þ xÞ½O þ ð1 þ xÞ=2 Xe This reaction results in the Sr2MnGaO5xF1þx oxyfluorides with x ¼ 0.22, 0.46 and 0.61. The Sr2MnGaO4.78F1.22 oxyfluoride has a tetragonal perovskite structure with a ap, c 2ap because the layered ordering of the Mn and Ga cations is preserved. In the Sr2MnGaO4.78F1.22 structure the MnO6 octahedra are characterized by two short apical ˚ and four long equatorial Mn-Oeq distances of 1.928 A ˚ , which Mn-Oap distances of 1.876 A was interpreted as an ‘apically compressed’ Jahn-Teller distortion, in contrast to the ‘apically elongated’ distortion in Sr2MnGaO5þ. It is interesting to compare the structures of oxygen-doped and fluorine-doped Mn-based brownmillerites with different Mn oxidation state (see Table 13.1). In oxygen-doped Sr2MnGaO5þ samples the reduction of Mn valence occurs concomitantly with an increasing amount of oxygen vacancies (decreasing ). In fluorinated Sr2MnGaO5xF1þx samples all anion positions are fully occupied, and the reduction occurs due to a replacement of oxygen by fluorine (increasing x). In both cases the Mn-Oeq bonds elongate with decreasing Mn oxidation state, following an increase of the Mn cation size upon reduction (Figure 13.23). On the other hand, Mn-Oap distances demonstrate different behaviour for the oxygenated and fluorinated samples. For oxygenated samples reduction leads to elongation of the Mn-Oap bonds, whereas the fluorinated samples show shortening of these bonds (Figure 13.23). In the oxygen-doped compounds the elongation of the Mn-Oap distance is to some degree compensated by shortening the Ga-Oap distance due to decreasing the Ga coordination number down to 4. For the fluorinated compounds the Ga-Oap distance does not alter since no changes in the coordination number of Ga occur. In the fluorinated sample the uncompensated apical tension of MnO6 octahedra would result in abnormally long Sr-O separations. One can speculate that in Sr2MnGaO5xF1þx the apically compressed Jahn-Teller distortion is more favourable for optimal Sr-O bonding. Sr2MnGaO4.78F1.22 is antiferromagnetically ordered below TN ¼ 70 K with the Mn magnetic moment aligned parallel to the (MnO2) plane [176,177]. The Mn spins are FM coupled between the planes, as it would be expected from ‘diagonal’ superexchange through completely filled Ga(O,F)2 layer, but in the plane the Mn spins form FM rods, which are AFM coupled to each other. In this structure one may propose that the z2 orbitals are orientationally ordered along the a-axis and the interaction along the b-axis is predominantly antiferromagnetic.

Table 13.1 Comparison of structural parameters and magnetic structures for Sr2MnGaO5.41 and Sr2MnGaO4.78F1.22 with nearly the same VMn

VMn d(Mn-Oeq) d(Mn-Oap) ﬀMn-Oeq-Mn CN(Ga) Magnetic structure

Sr2MnGaO5.41

Sr2MnGaO4.78F1.22

þ3.82 ˚ 1.904 A ˚ 1.985 A 180 4–5 FM coupled AFM layers

þ3.78 ˚ 1.928 A ˚ 1.876 A 180 6 AFM coupled FM rods

Properties of Cu- and Mn-based Oxyfluorides

415

2.50

2.40

d(Mn-O)ap (O)

2.30

d, Å

2.20

2.10

2.00

d(Mn-O)eq (O)

d(Mn-O)eq (F)

1.90 d(Mn-O)ap (F)

1.80 3.00

3.20

3.40 VMn

3.60

3.80

4.00

Figure 13.23 Interatomic distances vs. the Mn oxidation state (VMn) dependences for Sr2MnGaO5þ (O) and Sr2MnGaO5xF1þx (F)

Comparing the oxidized Sr2MnGaO5.41 and fluorinated Sr2MnGaO4.78F1.22 compounds one can see that the coordination environment of Mn cations, orbital ordering type and magnetic structure depend on the quantity of anions in the magnetically inactive Ga layers, but at the same time the oxidation state of Mn is almost the same due to differences in the formal charges of oxygen and fluorine. The structures of both compounds are far from the initial brownmillerite structure. They are layered perovskites with ordering of Mn and Ga cations. These ordered perovskites could not be prepared by direct solid state reaction and were obtained as a result of oxidation, reduction and fluorination reactions performed at soft conditions, preserving the layered cation arrangement of brownmillerite.

13.5

Conclusions

Numerous studies on fluorinated cuprates and manganites demonstrate that fluorination of the preformed complex oxides can effectively be used to adjust their properties for electronic applications. By fluorination, structural transformations can be induced, leading to compounds that cannot be obtained by a direct high-temperature solid state reaction. The directions of structural changes under fluorination depend on the nature of the fluorination reaction (anion insertion or replacement), the type of the crystallographic

416

Functionalized Inorganic Fluorides

positions available for fluorine (anion vacancies or interstitial sites) and the amount of extra anions inserted. Some general trends can be summarized: 1. A replacement of oxygen anions by an equivalent amount of fluorine anions causes a slight variation of lattice parameters due to differences in ionic radii of O2 and F and elongation of the metal-oxygen bond lengths due to reduction of the transition metal cations. 2. Filling anion vacancies with inserted fluorine anions completes the coordination polyhedron of the transition metal cation to an octahedron, which in the case of Jahn-Teller active cations leads to anisotropic coordination with four short equatorial and two longer apical metal-oxygen bonds. 3. Insertion of fluorine anions into the oxide structure with ordered anion vacancies suppresses the structural distortions caused by ordering of the oxygen anions and vacancies. 4. Insertion of fluorine anions into interstitial positions forces structural rearrangement in order to diminish the repulsive forces between the interstitial anions and anions in neighbouring layers. Fluorination of complex oxides was used as a tool to adjust the charge carrier concentration, which is required for the appearance of superconducting properties in cuprates. The fluorinated cuprates often demonstrate Tc not lower, and even higher than their oxide prototypes (see Table 13.2). In manganites, along with changing the Mn oxidation state, fluorination alters possible superexchange paths and orbital ordering patterns, affecting the magnetic properties and the magnetic structures. Additional information on the structure and properties of fluorinated cuprates and manganites, as well as on other complex transition metal oxides, can be found in several reviews [37,138,178–182]. Table 13.2

Superconducting complex copper oxyfluorides

Oxyfluoride

Rc (K)

Oxygenated prototype

Rc (K)

YBa2Cu3O6F2 Y2Ba4Cu7O14F2 HgBa2CuO4F0.24 HgBa2CaCu2O6Fd Hg0.8Ba2Ca2Cu3.2O8Fd Sr2CuO2F2þd La2CuO4Fd? d £ 0 18 Nd2CuO3.7F0.3 PbY0.8Ca0.2Ba2Cu3O7F2 Sr2CaCu2O4.6F2 Sr2Ca2Cu3O6.2F3.2 Sr2Nd0.2Ca0.8Cu2O5F

94 62 97 128 138 46 35–40 27 50 99 111 85

YBa2Cu3O6.95 Y2Ba4Cu7O14.92 HgBa2CuO4.12 HgBa2CaCu2O6.22 Hg0.8Ba2Ca2Cu3.2O8þd Sr2CuO3 La2CuO4.032 Nd2xCexCuO4 PbY1xCaxBa2Cu3O7þd – – –

92 80 97 127 134 – 38 24 90 – – –

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14 Doping Influence on the Defect Structure and Ionic Conductivity of Fluorine-containing Phases Elena I. Ardashnikova, Vladimir A. Prituzhalov and Ilya B. Kutsenok Department of Chemistry, Moscow State University, Moscow 119992, Russia

14.1

Introduction

There are a few requirements that ensure high fluoride-ion mobility in solids. • • • • •

High concentration of defects (ion vacancies or interstitials) in the structure. Various structure positions of fluoride-ions should have equal or similar energy states. Low cation-fluorine bond energy. Free fluoride-ions diffusion paths in the whole crystal lattice. Presence of highly polarizable i.e. light deforming cations which promote anion diffusion. • Local crystal structure distortions caused by the domain structure and micro-inclusions. The compounds with fluorite- and tysonite-type structures meet all these requirements very well (Table 14.1). The high ionic conductivity of these compounds combines with relative low electronic conductivity (usually no more than 0.1 % of the total conductivity [1]). This is a significant reason for their application as solid electrolytes.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

424

Functionalized Inorganic Fluorides

Table 14.1 Conductivity of some fluoride-ion conductive solid electrolytes with fluorite- and tysonite-like structures [2] (references to original papers are in Reprinted with permission from [2]). Copyright (2007) Pleiades Publishing Inc. FSE

T, K

Conductivity

electronic, S/cm CaF2

873

Sr0.8La0.2F2.2 Ba0.95La0.05F2.05 Ba0.6La0.4F2.4

575 667 518

PbF2 PbF2 PbF2 PbF2 PbF2 Pb0.99Y0.01F2.01 Pb0.66Y034F2 LaF3

545 581 293 373 325 423 765 773 293 373 543 297

378 LaF3 293 R1yMyF3y R ¼ La, Ce, Pr, Nd M¼Ca, Sr, Ba, Eu2þ 297 La0.95Sr0.05F2.95 La0.94Ba0.06F2.94 La0.9Ba0.1F2.9 Ce0.95Ca0.05F2.95

293 393 573

e ¼ 2 106 h ¼ ð1 5Þ 108 el ¼ 2 105 el ¼ 5 105 el ¼ 2:5 105 el ¼ 4:5 105 el ¼ 8 105 e ¼ 7 1010 e ¼ 3:5 107 e ¼ 5 109 h ¼ ð1 6Þ 1011 e ¼ 3 106 ; h ¼1 108 e ¼ 6 109 ; h ¼7 105 e ¼ 5:5 103 e ¼ 3:7 108 e ¼ 1:2 106 e ¼ 5 107 h ¼ 5 108 ðp F 2 ¼105 PaÞ

ionic, S/cm

8 102 8 102 6 102 1 101 2 101 1.8 108 1 105

9.2 101 4 107 1.4 105 3 102 1 106

5 e ¼ 1 107 ; h ¼5:5 1011 2.7 10 5 9 3 10 5 104 e < 10

h ¼ 1 106 ðpF2 ¼105 PaÞ e ¼ 1 106 el ¼ 5 109 e ¼3 107 el ¼3 107

Type of electrochemical cell

1 104 4 105 2 104 2 102

M, MFn|FSE|C M ¼ Ni, Co, Y Ni, NiF2|FSE|C Ni, NiF2|FSE|C Ni, NiF2|FSE|C Ni, NiF2|FSE|Ag Pb|FSE|C

Pb|FSE|Au Au|FSE|Ni, NiF2 Cu, CuF2|FSE|Au Pb(l)|FSE|C Pb|FSE|C Sn|FSE|C Pt|FSE|Pt, F2 pF2 ¼1033 105Pa La|FSE|C Ag, AgCl| KCl(l)KF(l)| FSE|M Pt|FSE|Pt, F2 pF2 ¼1033 105Pa La|FSE|RuO2 La|FSE|Pt C, Pb(l)|FSE|C

Note: el is the total electronic conductivity and e and h are the free-electron and hole conductivities, respectively.

Temperature dependences of conductivity of various fluoride electrolytes in comparison with oxide ones are presented in Figure 14.1 [3]. Fluoride phases with various compositions could be used as ionic conductors at different temperatures from 300 K. Unlike most oxide conductors, which could be used as solid electrolytes at certain pressures, only fluoride conductors keep their ionic conductivity in a wide range of fluorine partial pressures, pF2 . The temperatures and partial pressures of oxygen for various oxides or of fluorine for some fluorides where these phases have mainly ionic conductivity are shown in Figure 14.2 [4]. There is almost no information about proton conductivity of fluorides in the literature. To avoid hydrolysis and pyrohydrolysis the synthesis of fluorides and their conductivity

Doping Influence on Fluorine-containing Phases Igσ, Om–1 cm–1 (ZrO2)0.85(Y2O3)0.15

425

β

–1

PbSnF4 LaF3 α

–3 K0.06Bi0.94F2.88 β-PbF2

CaF2 –5

LaF3 Sr0.6La0.4F2.4

PbFCI BaF2 1.0

2.0

3.0

103/T, K–1

Figure 14.1 Temperature dependences of conductivity of some fluoride and oxide single crystals. (Reprinted with permission from [3] Copyright (1995) Kristallographia Publishing Company.)

400

600

2000 1600 1400 1200 1000 900 800

t, °C

45 25

ZrO2 (CaO) ThO2 (Y2O3)

lg px 2, Pa

5 –15

ThO2 ZrO2 (CaO)

ThO2 (Y2O3)

ThO2

–35

Th, ThF4

Ca, CaF2

–55

CaF2 –75

0

5

10

15

1/T · 103, K–1

Figure 14.2 Ionic conductivity regions on the log px2¼ f(1/T) (X ¼ O,F) diagram for some oxide and fluoride phases. (Reprinted with permission from [4] Copyright (2003) IKC Akademkniga.)

426

Functionalized Inorganic Fluorides

measurements is usually carried out in vacuum or in a dry gas atmosphere (argon, air, oxygen). Otherwise proton conductivity can arise, especially in the hydroscopic nonstoichiometric phases, for example KxBi1xF32x solid solutions, because of uncontrolled pyrohydrolysis of such phases and consequently F ions replacement by OH groups having almost similar size. In Figure 14.3 the conductivity of various fluoride-ion solid electrolytes with tysoniteand fluorite-like structures is presented as a function of the ionic transfer activation energy (Ea). It can be seen that a suitable fluoride solid electrolyte with optimal properties for any applications could be chosen or its targeted search could be carried out. Igσ500K, S/cm Ag4Rbl5 (Ag+)

0

1

β-Al2O3 (Na+)

2

–2 –

LaF3: Eu2+ (F )

–4

–6 Zr0.84Y0.16O1.84 (O2–)

–8

–10

–12 0

0,5

1

1,5

Ea, eV

Figure 14.3 Conductivity of various fluorine and oxygen conductors as a function of the ionic transfer activation energy. (Reprinted with permission from [2]. (1–fluorite-like phases, 2–tysonite-like phases). Copyright (2007) Pleiades Publishing Inc.)

Recently a lot of reviews on ionic conductivity of inorganic fluorides were published. A short history of these investigations starting with Faraday’s basic work of 1834 up to their current applications is presented in [2]. The structure and ionic conductivity of Pb1xAlxF2þx, M1x(U or Th)xF2þ2x (M ¼ Ca, Sr, Ba, Pb) fluorites and CeF3, Ce1yCdyF3y tysonites are described in [5]. In a detailed review [6], various fluoride-conductive electrolytes: perovskites MPbF3 (M ¼ K, Rb, Cs), B-deficit perovskites – tysonites

Doping Influence on Fluorine-containing Phases

427

(LaF3-type structure), double perovskites (Na3AlF6-type structure) and fluorites and their derivatives: ß-PbF2 and nonstoichiometric PbF2-based phases (containing K, Bi, Y, Al, Th, Zr, Ta), PbSnF4, Ca1xYxF2þx, M1xBixF1þ2x (M ¼ K, Rb) are discussed. That is why in the current chapter we will confine ourselves to the ionic transfer discussion of fluorine-conductive fluorites and tysonites from a slightly different standpoint; we will try to observe an influence of various doping types on the structure features and the defect structure of nonstoichiometric fluorine-containing phases. It is well known that the conductivities of single crystals, ceramic and thin film samples are slightly different. Growing single fluoride crystals is a special and rather complicated task, so the results of studying ceramic materials by different authors conducted in identical conditions will be discussed in the present review. Nevertheless, comparison of conductivity values measured in different laboratories is not always correct; that is why sometimes only a general tendency observed by different authors is important. At the same time, the macrostructure of samples is of great importance and is affected by the synthesis history and therefore by the nonequilibrium degree of the sample as well as by the presence of ordered phases in the corresponding systems.

14.2

Influence of Oxygen Ions on Fluoride Properties

The partial replacement of fluoride ions by oxygen ones in crystalline fluorides is important. Even low oxygen content in fluorides changes their chemical and physical properties markedly. For example, we can cite conductivities, melting points and phase transformation temperatures of rare-earth (RE) metal trifluorides [7]. It also stabilizes nonstable modifications, for example the high-pressure modification of BiF3 [8]). Sometimes in the process of searching for new phases scientists mistakenly construe oxyfluoride phases as pure fluoride ones (see for example [9]). 14.2.1

Pyrohydrolysis

Pyrohydrolysis is a substantial but often uncontrolled factor in the production process and measurements although it strongly influences the ionic conductivity of fluorides. Pyrohydrolysis is the reaction between fluorides and the water adsorbed or existing in the surrounding gas phase at high temperatures (to). Unlike hydrolysis that takes place at low temperatures and leads to F -ions replacing by OH-ions, in process of pyrohydrolysis F -ions are replaced by O2 -ions. Thus, in the process of pyrohydrolysis, oxygen ions enter into the fluoride structure according to the following equation: to

2 F þ H2 O ¼ O2 þ 2 HFðgÞ " For example, BiF3 undergoes many changes in process of its pyrohydrolysis; a few oxyfluorides and even pure oxides could arise according to the below scheme: to

to

to

to

to

to

H2 O

H2 O

H2 O

H2 O

H2 O

H2 O

BiF3 ! -BiOx F3 2x ! -BiOx F3 2x ! BiOF ! -BiOx F3 2x ! Bi7 O9 F3 ! Bi2 O3

428

14.2.2

Functionalized Inorganic Fluorides

Heterovalent Oxygen Substitution for Fluoride Ions

Heterovalent oxygen substitution for fluoride ions occurs relatively easily because both ions are rather similar. They have the same electronic structure, similar weights and sizes and other characteristics. Of course there are a few differences between them: the charge, electronegativity and donor activity. A small distinction in the anion-cation bond energy leads to the formation of oxyfluorides with a disordered anion sublattice. Nevertheless there is a preferable coordination of F or O2 ions around some cations. To maintain the electroneutrality of the structure two fluoride ions should be replaced by an oxygen ion: 2F $ O2. This replacement results in decreasing fluoride-ion interstitials or increasing anion vacancies concentration. That is why the oxygen introduction sometimes leads to decreasing conductivity as in Pb0.99K0.01F1.99 [10] and sometimes to its increasing as in Sr0.9Lu0.1F2.1 (Figure 14.4) [1].

In σT (Sm–1K)

2

1 2

–3

–8 1.0

1.5

2.0

103/T (K–1)

Figure 14.4 Temperature dependence of the Sr0.9Lu0.1F2.1 conductivity: 1- in the air, 2- vacuum 102 Pa. (Reprinted with permission from [1] Copyright (1989) Elsevier Ltd.)

It was shown for Bi1xBax(O,F)3 oxyfluorides with tysonite structure that introducing oxygen ions into modification I of this phase leads to the conductivity decrease and into modification II to its increase (Figure 14.5) [11]. The results of oxygen insertion depend on several factors such as the structure of nonstoichiometric phases, the oxygen content and the oxygen positions in the structure as well as the types and concentrations of other defects of the crystal lattice. Thus the introduction of oxygen into the fluoride structure can be considered as an effective tool for the ionic transfer control and its investigation.

Doping Influence on Fluorine-containing Phases σ*10–5

429

Bi0,96 Ba0,04 Oy F2,96-2y

Om–1 cm–1

4

3

2

1

0 0,04

0,08 y, mol. %

0,12

0,16

Figure 14.5 Dependence of conductivity at 300 K as a function of oxygen content for two modifications of Bi0.96Ba0.04OyF2.962y tysonites [11]

14.2.3

Ionic Conductivity of Oxyfluorides

In this review we will consider anion-conductive fluorides and oxyfluorides because the occurrence of oxygen in these phases is an additional parameter that controls the defect structure of fluorite- and tysonite-like solid solutions by increasing the anion vacancies or decreasing the interstitial ions. Fluoride ion replacement by oxygen in some oxyfluoride conductors, for example in Bi1xBax(O,F)3 tysonite-like (Figure 14.5) [11] or Bi1xTex(O,F)2þ fluorite-like phases (Figure 14.6) [12] leads, together with the effects mentioned above, to a decrease in the mobile fluoride ion concentration. A similar conductivity dependence – the conductivity raises with the fluoride ion concentration – was also found for the fluoride-conductive cubic LaOF phase (Figure 14.7) [13]. The type of charge carriers in this case was confirmed by means of the Tubandt method. Using Ni þ NiO/LaOF/Ni and Ni þ NiF2/LaOF/Ni galvanic cells, it was discovered that the high temperature LaOF phase does not have any significant oxygen conductivity at the temperatures under investigation (the values of the ionic transfer numbers are tF ¼ 0.989–1.001, tO < 102) [13]. There are different points of view on the kind of charge carriers responsible for the conductivity in oxyfluoride phases. For example Ln2 Ln 02 O3 F6 phases (La2Eu2O3F6: ¼ 0.8 Sm1, tO ¼ 0.7 at 773 K and Nd2Eu2O3F6: ¼ 2.0 Sm1, tO ¼ 0.9 at 773 K) were considered as oxygen-conductive phases [14,15] (Figure 14.8). Their conductivities do not depend on the oxygen pressure. To confirm the charge carrier type, electrolysis of Nd2Eu2O3F6 was carried out (973 J, 12 hours, 1 mA, quartz ampoule) with a Ni þ NiO anode [16]. After the electrolysis the only phase at the anode was NiO and no trace of fluoride was detected.

430

Functionalized Inorganic Fluorides 1

IgσT, Om–1 cm–1 K

Bi0.7 Te0.3 O1.3–y F0.7+2y

0

-y=0 - y = 0.1 - y = 0.25 - y = 0.4

–1

–2 –3 –4 1,6

2,0

2,8

2,4

3,2

1000/T, K–1

Figure 14.6 Temperature dependence of conductivity of the fluorite-like phase Bi1xTex(O,F)2þd (x ¼ 0.3) [12]

0.0

log (σ/S m–1)

LaO1.01F0.98 –0.2

LaO1.00F1.00 LaO0.99F1.02

–0.4

LaO0.98F1.04

–0.6 –0.8 –1.0 –1.2 –1.4 7

8

9

10

[1/(T/K )] × 104

Figure 14.7 Conductivity of the LaO1xF1þ2x cubic fluorite-like phases. (Reprinted with permission from [13] Copyright (2006) Wiley-VCH.)

Doping Influence on Fluorine-containing Phases

431

Rare-earth fluoride, Ln’F3 La Ce

Pr Nd Sm Eu Gd Tb Dy Ho

Y

Er Tm Yb Lu

La CeO2

Rare-earth oxide, Ln2O3

Pr6O11 Nd Sm Eu Gd Tb4O7 Dy Ho Y Er Tm Yb Lu

Figure 14.8 Electrical conductivity of binary RE oxyfluorides Ln2Ln’2O3F6 measured at 973 K under oxygen partial pressure of 0.13 Pa. , more than 1S m1; C, 0.1 1S m1; , less than 0.1S m1. (Reprinted with permission from [15] Copyright (2000) Elsevier Ltd.)

•

Because the easy oxygen introduction into fluorides changes essentially their conductivity some of such phases are rather sensitive to the gas atmosphere composition and could have different applications. For example LaF3 and PbSnF4 were used as oxygen and moisture sensors [17–19].

14.3

Cation Doping of Fluorides

Cation doping of fluorides always leads to a crystal structure distortion. The degree of distortion depends on the size ratio and properties (in the first place, electronic structure and electronegativity) of dopant and matrix cations as well as on the replacement type: isovalent or heterovalent replacement. 14.3.1

Isovalent Replacement in the Cation Sublattice

Isovalent replacement in the cation sublattice has not been studied precisely, probably because it does not lead to the formation of additional anion defects that increase the ionic conductivity. Nevertheless the isovalent replacement in the cation sublattice of fluorite (for example PbF2 doped by CaF2 or CdF2 [10]) or tysonite (for example LaF3

432

Functionalized Inorganic Fluorides

doped by NdF3 [20]) results in the conductivity increase. For example, the conductivity of Pb0.67Cd0.33F2 at 473 K is 50 times higher than that of PbF2 and the conductivity of Nd0.67La0.33F3 at 293 K is two times higher than that of LaF3. According to [2,21] these facts are due to changes in the bond ionicity degree and solid solutions properties (melting points, unit cell dimensions, energy of formation and migration of mobile defects etc.). 14.3.2

Heterovalent Replacement in the Cation Sublattice

Heterovalent replacement in the cation sublattice leads to an increase in the number of anion defects, for example vacancies (R3þ $ M2þ þ VF) in R1xMxF3x solid solutions or interstitial ions (M2þ $ R3þ þ Fi) in M1xRxF2þx solids solutions. Of course, this replacement greatly increases the conductivity of doped phases and will be discussed in detail hereinafter. Very wide homogeneity regions of the anion-excess fluorite-like M1xRxF2þx (0 < t < 0.5) solid solutions [7] are another unusual feature that will also be considered. For the anion-deficit tysonite-type Bi1xMx(O,F)3 (M ¼ Na, Sr, Ba) solid solutions a rather strange influence of the dimensional factor on the homogeneity region width was detected. This was an increase of M cations content, which was found from x ¼ 0.08 (in systems with M ¼ Na, Sr having close cation radii to the matrix cation radius) to x ¼ 0.17 (for the large M ¼ Ba cation) at 873 K [11,22]. Possibly, the homogeneity region widening with an increase of the dopant ionic radius could be explained by two factors compensating for each other – the large dopant cation and the presence of anion vacancies. The first factor leads to the unit cell expansion and the second one leads to its reduction because the larger the unit cell, the higher its capacity to adopt anion vacancies. However, in the R1xMxF3x and M1xRxF2þx examples mentioned above, the conductivity is affected, in addition to an increase in anion defect numbers, by two other factors: the size factor (due to cation replacement and anion sublattice defects) and the change of the electronic structure (due to the introduction a dopant cation). The effect of the size factor for similar solid solutions is quite predictable: the conductivity slightly increases and the activation energy Ea slightly decreases with the raise of the unit cell parameters (and correspondingly with an expansion of conductivity channels). The effect of the electronic structure of the dopant cation, for example the availability of a lone electron pair is more significant.

14.4

Active Lone Electron Pair of Cations and Ionic Conductivity

A lone electron pair deforming and sometimes changing the crystal structure and conductivity channels influences the conductive properties of materials. There is a good example of the lone electron pair’s influence, namely the oxyfluoride solid solutions based on NdOF and BiOF having a fluorite-like structure. Both phases have an identical ˚, anionic composition, similar structures and close cation radii (RBi3þ ¼ 1.31 A ˚ [23]). RNd3þ ¼ 1.25 A

Doping Influence on Fluorine-containing Phases

433

The conductivity of BiOF [24] is higher than that of NdOF [25] and the doping increases anionic conductivity of both phases [26] (Table 14.2, Figure 14.9). Obviously, it is caused by the easy deformability of the bismuth cation distorting the crystal lattice of the oxyfluoride. Actually, at the same fluoride ion content, the conductivity of the phase Bi1tTet(O,F)2þ containing two easy deformable cations (Bi3þ and Te4þ) with lone electron pairs and having a more distorted fluorite lattice is higher than the conductivity of Nd1xTex(O,F)2 (Table 14.2]. At the same time the conductivity activation energies are quite close (Ea ¼ 0.40–0.50 eV) for both phases. The conductivity mechanism is probably similar in these phases [12]. IgσT, Om–1 sm–1 K

2

Bi0.90 Te0.10 O1.10 F0.90

1 0

Nd0.9 Te0.1 O1.2 F0.7

–1 –2 NdOF

–3

Bi0.95 Y0.05 OF

BiOF

–4 1,0

1,5

2,0 1000/T, K–1

2,5

3,0

Figure 14.9 Ionic conductivity of the fluorite-like phases based on MOF (M ¼ Bi, Nd) [after 12, 24–26]

Table 14.2 Nd)

Ionic conductivity and the structure of fluorite-like phases based on MOF (M ¼ Bi,

Phase [References]

Crystal System Relation to the fluorite lattice ao p (layer structure) BiOF [24] tetragonal at ¼ ao p2/2 cp p Bi0.95Y0.05OF [24] trigonal ah ¼ (ao p2/2) 3, ch ¼pao 3 Bi0.9Te0.1O1.1F0.9 [26] trigonal ah ¼ 3ap o 2/2, ch ¼ aop 3 NdOF [25] trigonal ah ¼ ao 2/2, ch ¼ 2ao 3 NdOF [25] Cubic ac ¼ ao ac ¼ ao Nd0.9Te0.1O1.1F0.9 [26] Cubic

400 K, O1 cm1 3.2106 1.3104 6.3104 107 at 723 K 4105at 823 K 2.5106

PbSnF4 is the most conductive phase among fluorides and both cations Pb2þ and Sn2þ have lone electron pairs. The PbSnF4 structure is strongly distorted in comparison with ß-PbF2 fluorite and Pb1xSnxF2 solid solutions [27] due to the ordered location of Pb and Sn cations. The layered structure PbPbSnSnPbPbSnSn causes the conductivity anisotropy to be smoothed in different-oriented pressed powder samples.

434

Functionalized Inorganic Fluorides

A number of closely related fluorite-like phases of PbSnF4 with orthorhombic, tetragonal and monoclinic structures were synthesized by different methods. The tetragonal ordered -PbSnF4 phase (P4/nmm) is stable at ambient temperatures and could be prepared by solid-state synthesis of constituent binary fluorides [6]. The structure of -PbSnF4 remains tetragonal at high temperatures [28]. The phase transitions ! ß ! occur at 608 and 672 J, respectively; its melting point is 693 K [29]. The lattice pﬃﬃﬃ parameters of -PbSnF4 are related to those of the parent fluorite structure by a ¼ afl/ 2, c ¼ 2afl. The stabilization of the layered structure is generally attributed to 5s2 electron lone-pairs of Sn2þ cations. They point to the interlayer space between the two Sn2þ sheets, creating cation-centred pseudo-octahedral SnF5E units (where E represents the lone pair) [29].

F(1) Sn F(2) F(4) Pb F(3)

Figure 14.10 Structure of a-PbSnF4 showing the cation layers and the four fluoride site; nominal vacant F(1) site; the mobile F(2) and F(4) sites between Sn and Pb layers; and static F(3) sites between two Pb layers. (Reprinted with permission from [30] Copyright (2008) American Chemical Society.)

It was shown by powder neutron diffraction, 19F field-cycling NMR and MD simulation that the static F(3) sites are located between two Pb layers in -PbSnF4 and the mobile F(2) and F(4) sites are between Sn and Pb layers [6,29,30]. The time-averaged anion density and two-dimensional diffusion within -PbSnF4 have recently been successfully reproduced within MD simulations, which essentially treat the ‘lone-pairs’ as an extreme manifestation of cation polarizability. As a consequence, the structural behaviour of -PbSnF4 can be interpreted in terms of a higher ‘effective’ polarizability of the smaller Sn2þ ion [6]. The authors of [30] recently proposed a detailed mechanism for conductivity of this compound.

Doping Influence on Fluorine-containing Phases

14.5

14.5.1

435

Peculiarities of the Defect Structure of Nonstoichiometric Fluorite-like Phases Fluorite Structure

Fluorite is a mineral composed of calcium fluoride, CaF2. This mineral gives its name to the fluorite structural type. Its structure (Figure 14.11a) has cubic symmetry, the space ˚ [31]). The structure can be described as a face-centred cubic group is Fm-3m (a ¼ 5.471 A (fcc) array of cations in which all the fourfold coordinated interstices are filled with anions and the sixfold coordinated ones are empty (Figure 14.11a). The coordination number of cations is eight (cube) and that of anions is four (tetrahedron). There are many compounds of various chemical classes that crystallize in this structural type: fluorides of alkaline-earth elements, lead and cadmium; high-temperature modifications of zirconium and hafnium oxides, solid solutions M1xRxF2þx (M ¼ Ca, Sr, Ba, Pb, Cd; R ¼ RE elements) [7,21, 31 and others] Ba1xBix(O,F)2þ [32], fluorite-like modifications of MOF (M ¼ RE elements, Bi) and M1xTex(O,F)2þ oxyfluoride phases [33,34], solid solutions in the BiOF-YOF system [24] etc. Structural investigations of Ca1xRxF2þx, Sr1xRxF2þx, and Ba1xRxF2þx solid solutions were carried out for all RE elements, some RE elements of both subgroups, and for RE elements of the cerium subgroup, respectively. The basic result of these researches was the establishment of the cluster structure of anion-excess fluorite-like solid solutions. The following methods were used in these studies – X-ray, electronic and neutron diffractions, neutron and proton diffusion scattering, EXAFS, optical, laser and luminescence spectroscopies, 19F NMR, EPR etc. Currently, associations of structural defects (clusters) are the basic structural units in nonstoichiomtric M2þ1xR3þxF2þx phases. Clusters formed by RE ions, interstitial anions and anion vacancies are located in the initial fluorite matrix. 14.5.2

Defect Clusters

Formation of an anion-excess solid solution can be described by the following quasichemical equation (by the example of MF2 doped by RF3): 0

RF3 ! R•M þ 2FFx þ Fi ! MF 2 where M is the bivalent cation of the ‘‘matrix’’, R is the trivalent doping cation, FF (or Fn) is a fluorine atom occupying the position (¼;¼;¼), and Fi is a fluorine atom occupying an interstitial position. Possible variants of the interstitial Fi ions location are presented in Figure 14.11. There are a few positions for interstitial anions Fi (Figure 14.11b): • F0 , fluorine atom displaced from the normal position in direction <011>. • F00 , fluorine atom displaced from the normal position in direction <111> • F000 , fluorine atom also displaced from the normal position in direction <111>. The last one is not a real interstitial position but just a ‘relaxed position’ – normal with very small displacement. There are different systems for marking interstitial positions in the scientific literature [21,35,36].

436

Functionalized Inorganic Fluorides

Ca2+ –

F

V

- cations - anionic vacancy

a)

- anion in the initial position - interstitial anion

b)

Figure 14.11 Fluorite structure: a) stoichiometric fluorite structure; b) defects of the anion fluorite sublattice

An interconnected anion displacement occurs in the process of MII1xRIIIxF2þx solid solution formation to avoid very short F-F interatomic distances, which would arise if interstitial fluorine atoms would be statistically distributed in the fluorite structure. Then some fluorine atoms have to shift from their initial crystallographic Fn positions. To compensate for the charge of interstitial anions, a part of M2þ cations should be replaced by R3þ. Additional fluorine atoms occupy interstitial positions only and M2þ and R3þ cations practically retain their positions in the fcc lattice. The fluorite structure allows MII1xRIIIxF2þx solid solutions to be obtained with high degrees of replacement up to 50 % of RF3. At low doping level (x £ 0.01) doublet pairs (R3þ-Fi) arise to maintain a net balance between positive and negative charges of the structure. The interstitial anion concentration increase (x > 0.01) leads to the formation of defect associates (clusters). Clusters can be denoted by a set of four numbers n1:n2:n3:n4 corresponding to the quantity of anionic vacancies and interstitial ions in F0 , F00 and F000 positions, respectively. The last number is often not presented since for main cluster types it is equal to 0. More than 30 types of various clusters were found out: 8:12:0; 8:12:1; 1:0:n (n ¼ 2, 3, 4); (n þ 1):2n:1; (2n þ 2):3n:2; 4:4:3, 3:2:3; 2:2:2, etc [7,37]. Several different types of clusters are often present in a phase. For instance, coexistence of 8:12:1 and 1:0:3 clusters for the Ca1xRxF2þx (R ¼ Er, Ho) solid solutions [38], combination of 8:12:1 and 4:4:n (where n is between 1 and 4) clusters for the Ca1xYxF2þx solid solutions [39,40], combination of cubooctahedral and 1:0:3 clusters for the Na0.5xBi0.5þx(O,F)2þ solid solutions [41], and combination of cubooctahedral and 1:0:4 clusters for the Cd1xRxF2þx (R ¼ La, Nd) solid solutions [21] were assumed. Because the cuboctahedral clusters are the most widespread and the largest scale ones, we will therefore consider them in detail. All eight normal anion positions (FF) in the fluorite unit cell containing such a cluster are vacant. Twelve atoms F0 form the cuboctahedron (8:12:0). The position in the centre of the unit cell could be occupied by one additional fluorine atom F00 (8:12:1). The anion cluster is surrounded by an octahedron of R3þ cations located in face centres (Figure 14.12). The coordination polyhedron of R3þ cations is a square antiprism (coordination number equal to 8).

Doping Influence on Fluorine-containing Phases

Figure 14.12

437

Anion cuboctahedron coordinated by cations

The supercluster model [42,43] as a development of the cluster model considers not only anion cuboctahedral clusters but surrounding cations as well. A supercluster is an assembly of all cationic polyhedra having joint faces with the central anionic cubooctahedron. A supercluster has the [M14F68–69] formula because 8 M2þ cations located in cube corners with their closest anion surroundings from the initial fluorite matrix are added to the central anion cuboctahedron surrounded by 6 R3þ cations. In this formula, F69 arises if the position in the centre of cuboctahedron is filled (i.e. 8:12:1). According to [7,43], all RE ions and all types of structural defects are gathered in superclusters, for example in the form of a ‘rare-earth’ supercluster [M8R6F68–69] (Figure 14.13).

Figure 14.13 [M8R6F6869] supercluster and distorted cationic polyhedra around the anionic cluster: white colored – M2þ polyhedra, grey colored - R3þ polyhedra. (Reprinted with permission from [7] Copyright (2001) Bovis P. Sobolev.)

438

Functionalized Inorganic Fluorides

In a later work [21] it was suggested to consider various types of superclusters [M14F68] and to name their extreme compositions [M8{R6F36}F32] as octacubic cluster and [R8{M6F36}F32] as inverted octacubic cluster (Figure 14.14). In the M1xRxF2þx solid solutions, the formation of supercluster with variable R3þ dopant content [M14nRnF68] is possible. Such clusters have a size of around 1 nm.

F(8c) [M6F32]

Fint(48i )

[M6F36] [R6F36]

[M 2+] R 3+

[M8{R6F36}F32] Octacubic cluster

R 3+ M 2+

[R8{M6F36}F32] Inverted octacubic cluster

Figure 14.14 Superclusters in the M1xRxF2þx solid solutions. (Reprinted with permission from [21] Copyright (2008) Pleiades Publishing Inc.)

The model of the supercluster formation from cuboctahedral anion clusters and their surroundings of cationic polyhedra is quite applicable to other types of anionic clusters. The formation of a nonstoichiometric fluorite-like solid solution and the coexistence of various cluster types (cuboctahedral and tetrahedral) in it is well illustrated in Figure 14.15)[7].

Doping Influence on Fluorine-containing Phases

Fluorite structure

CaF2 M - Ca, Sr, Ba, Cd, Pb

Undefected structural block [M6F32]

Undefected structural block [M4F23]

Fc

Fc

Cluster of defects [R4F26]

[F8] Cluster of defects [R6F36] 2+2 Fi

439

Ff VFc

Fi

[F12]

Nonstoichiometric crystal Supercluster {M8[R6F68–69]}

Figure 14.15 Scheme of the nonstoichiometric M1xRxF2þx solid solution formation from two types of superclusters. (Reprinted with permission from [7] Copyright (2001) Bovis P. Sobolev.)

14.5.3

Ordered Fluorite-like Phases

The fact that almost all R3þ ions are concentrated in superclusters allows the crystal matrix to keep its structure similar to the initial one in a wide concentration interval.

440

Functionalized Inorganic Fluorides

With increasing x in M1xRxF2þx the formation of ordered phases often occurs. For example, in the BaF2-BiF3 system, cubic fluorite-like solid solutions were obtained up to x ¼ 0.45 at 973 K [44] or up to x ¼ 0.27 at 873 K [32]. For higher RF3 concentrations the formation of the fluorite-like ordered hexagonal Ba4Bi3F17 phase begins. In the BaF2-BiF3-Bi2O3 system, a more ordered fluorite-like phase, Ba2.1Bi0.9(O,F)6.8, was detected [32]. This phase could be obtained in presence of oxygen ions only. Information on some ordered fluorine-containing fluorite-like phases is presented in Table 14.3. In all these phases, the cation position is split. It means that both M2þ and R3þ cations are ordered in the crystal lattice only in classical cation positions of the fluorite structure. The main difference between initial and ordered phases consists in nonequivalence of anions packages in the crystal structure. Practically all the structures of the ordered phases studied by now, in fluoride systems with RE elements, belong to the class of phases with ordered cuboctahedral clusters. The position in the centre of cuboctahedron can be occupied by an anion (8:12:1) or can be vacant (8:12:0). In oxyfluoride ordered phases, the position in the centre of cuboctahedron is usually vacant, as for instance in PbZr3F6O4 [60], Pb8Y6F32O [61], Ba2.1Bi0.9(O,F)6.8 [59]. Table 14.3 Compositions and unit cell parameters for some ordered fluorine-containing fluorite-like phases Formula

Crystal System

MIR2F7

monoclinic

MIR2F7

monoclinic

MI3R5F18 MI7R13F46 MIR3F10 MII0,8R0,2F2,2 MII14R5F43 MII2RF7 MII2RF7

tetragonal orthorhombic cubic cubic hexagonal tetragonal tetragonal

MII9R5F33 MII17R10F64 MII5R3F16 MII8R5F31 MI7R6F31 MII4R3F17 MII4R3F17 MII4R3F17 R7X16 (X ¼ F þ O) R2X5 (X ¼ F þ O) MII2RX7 (X ¼ F þ O)

hexagonal cubic tetragonal hexagonal hexagonal hexagonal hexagonal tetragonal monoclinic hexagonal tetragonal

Space group

Cm

R-3 I4/m

Pa3

R-3 R-3 R-3

I4/m

Parameters pﬃﬃﬃ pﬃﬃﬃ a ¼ afl p 6,ﬃﬃﬃﬃﬃﬃﬃﬃ b ¼ afl 2, 9 c ¼ 3ap fl ﬃﬃﬃ =2 pﬃﬃﬃ a ¼ aflp6ﬃﬃﬃﬃﬃﬃﬃﬃ , b ¼ afl 2, c ¼ afl 9 =2 a ¼ 2afl a ¼ aflpﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ pﬃﬃﬃ a ¼ afl 19 =2 , c ¼ afl 3 pﬃﬃﬃﬃﬃﬃﬃﬃ 5 = , c ¼ 3a a ¼ afl . 2 fl pﬃﬃ 2 , c ¼ 3afl a ¼ afl 2 pﬃﬃﬃﬃﬃﬃﬃﬃ pﬃﬃﬃ a ¼ afl 7 =2 , c ¼ 2afl 3 a ¼ 3afl pﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ pﬃﬃﬃ a ¼ afl 13 =2 , c ¼ afl 3 pﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ pﬃﬃﬃ a ¼ afl 13 =2 , c ¼ afl 3 pﬃﬃﬃﬃﬃﬃﬃﬃ pﬃﬃﬃ 7 = , c ¼ 2a a ¼ aflpﬃﬃﬃﬃﬃﬃﬃﬃ 2 fl 3 pﬃﬃﬃ a ¼ afl pﬃﬃ7ﬃ=2 , c ¼ 2afl 3 / ﬃﬃ2ﬃ :, c ¼ afl a ¼ aflp pﬃﬃﬃﬃﬃﬃﬃﬃ 5= a ¼ aflp5 c¼ 2 ﬃﬃﬃ, b ¼ afl, p ﬃﬃﬃ afl a ¼ aflp2ﬃﬃﬃﬃﬃﬃﬃﬃ , c ¼ 2afl 3 a ¼ afl 5 =2 , c ¼ 3afl

R

MI, Ref. MII

Er

K

[45]

Ho

K

[42]

Yb Yb Y Yb Yb Yb Yb

Na Na K Ba Ca Ca Ba

[7] [7] [46] [47] [48] [49] [50]

Y Ca Yb Ca Er Ba Ho Ca Zr Na Yb Ba Bi Ba Gd Ba Bi7F11O5 Bi(O,F)2,45 Ba2.1Bi0.9 (O,F)6.8

[51] [42] [52] [53] [42] [54] [55] [56] [57] [58] [59]

Doping Influence on Fluorine-containing Phases

14.5.4

441

Phase Diagrams

Solid solutions usable as solid electrolytes are usually nonequilibrium since they are either quenched high-temperature phases or single crystals grown up from the melt. Nevertheless they are rather stable kinetically and show reproduced results for a long time. It is important to know what can occur after their long usage as electrolytes: whether phase transformations (for example at temperatures lower than the synthesis temperature) are possible according to the equilibrium phase diagram or whether crystallization of ordered phases can occur. Thus the investigation of the corresponding phase diagrams is very important for the further applications of nonstoichiometric solid solutions. Examples of some phase diagrams are presented in Figure 14.16 (isothermal section of a ternary phase diagram) and Figure 14.17 (binary phase diagrams).

- monophase samples - double-phase samples

BaF2

- ternary-phase samples Φ BiOF + Φ

R R+Φ+β

BiOF + Φ + β

R+β

1/2(Bi2O3)

BiOF

Φ+β

I

l+β+R

β

II III

BiF3

Figure 14.16 Solid solution in the BiF3–Bi2O3–BaF2 system at 873 K [34].R: Ba4Bi3F17 based phase, : -BiOyF32y based phase, F: fluorite-like BaF2 based solid solution, I, II: tysonite-like solid solutions, III: ordered tysonite-like phase BiOyF32y(0.13 < y < 0.23)

14.6

Ionic Transfer in Fluorite-like Phases

The conductivity of ionic crystals can be assumed to be due to the mobilities of the individual ions and electrons: ¼ q nmob mmob In this formula q, nmob and mmob denote respectively, the charge, the concentration and the mobility of charge carriers. In fluorite-like phases these charge carriers are mainly fluoride ions.

442

Functionalized Inorganic Fluorides

Figure 14.17 Phase diagrams of the SrF2–RF3 systems. (Reprinted with permission from [7,62] Copyright (1981) Elsevier Ltd.)

Doping Influence on Fluorine-containing Phases

443

The ionic conductivity of fluorite-like phases is connected to two types of anionic sublattice structural defects: excess interstitial fluoride ions Fi and vacancies in regular positions VF. The high defect concentration increases the conductivity. The introduction of excess fluoride ions into fluorite matrix by means of the heterovalent replacement leads to an increase in both the concentration and the mobility of charge carriers. As a result the conductivity can be increased by several orders of magnitude (Figure 14.18[63]). In this figure and further in this section, MF2–RF3 solid solutions will be considered as examples. As shown in Figure 14.18, the conductivity not only depends on the RF3 concentration, but also on the size factor, namely on the M2þ and R3þ ionic radii correlation since both of these factors define a defect cluster type existing in the solid solution structure. For example, the value of ionic conductivity of the nonstoichiometric Pb0.9R0.1F2.1 phase depends on the cation size ratio for different R (Figure 14.19) [10]. –logσ [S cm–1] 0

–logσ [S cm–1]

(a)

0

4

8

(b)

4

1

2.0

4

103/T,

K–1

8

8 8 5 6

3 3.2

(c)

0

4

2 12 0.8

–logσ [S cm–1]

12 0.6

2.2 103/T,

10

3.8 K–1

11

9

7 12 0.6

2.6 103/T,

12 4.6

K–1

Figure 14.18 Temperature dependences of conductivity for some fluorite-like solid solutions. (Reprinted with permission from [63]: a) based on CaF2 1–Ca0.999Gd0.001F2.001, 2–Ca0.8Lu0.2F2.2, 3–Ca0.8Gd0.2F2.2, 4–Ca0.7Gd0.3F2.3, b) based on SrF2 5–Sr0.998La0.002F2.002, 6–Sr0.8Lu0.2F2.2, 7–Sr0.79Gd0.21F2.21, 8–Sr0.6La0.4F2.4, c) based on BaF2 9–Ba0.998La0.002F2.002, 10– Ba0.7Tm0.3F2.3, 11–Ba0.8Gd0.2F2.2, 12–Ba0.5La0.5F2.5. Copyright (2006) Pleiades Publishing Inc.)

14.6.1

Defect Region Model

Features of the ionic transport in the fluorite-like solid solutions are well described by means of the ‘defect regions’ model [64,65]. The basic structural motive of such region is a defect cluster and a distorted fluorite matrix around it. Charge carriers in such a structure are fluoride ions in interstitial positions but only those that are placed on the periphery of the defect regions. The fluoride ions forming the defect cluster do not participate in the ionic transfer. The presence of free mobile fluoride ions around the cluster is a result of compensation of the defect region positive charge by means of additional fluoride ions, similarly to the positive charge compensation in the cuboctahedral cluster formation: [M6F32]20 þ 6 Fn ! [R6F37]19 þ Fi [66].

Functionalized Inorganic Fluorides

444

lg σ293 K (S/cm) Al

Lu Yb Er Y Dy Td Gd Sm

Sb

Ga

Sc In

La

–4

–5

–6 0.6

0.8

1.0

ri , Å

1.2

Figure 14.19 Ionic conductivity dependence of Pb0.9R0.1F2.1 phases, at 293 K, on the R3þ ionic radius. (Reprinted with permission from [10] Copyright (1997) Pleiades Publishing Inc.)

It is shown in [67] for Cd0.9R0.1F2.1 (R ¼ La-Lu, Y) that ionic conductivity (500K) of fluorites of various RE elements correlates with concentration of interstitial fluoride ions (Fi(48g) þ Fi(32f)1 þ Fi(32f)2) located on the periphery of clusters. These interstitial fluoride ion concentrations were measured by means of single-crystal X-ray diffraction (XRD). The concentration and mobility of charge carriers depend on the cluster type forming the defect region. At a RF3 concentration called the percolation limit – xper – defect regions combine themselves in joint conductivity channels. It results in a sharp conductivity increase in fluorite samples (Figure 14.20 [63]). –logσ [S cm–1] 4

Ea, eV

(a)

2

(b)

1.4

1 1.2

6 3

1.0 8

3 1

0.8

2

10 0

0.1

0.2

0.3

x

0.6

0

0.1

0.2

0.3

x

Figure 14.20 Concentration dependences of the conductivity (log ) at 500 K and of the conductivity activation energy for Ca1xRxF2þx solid solutions. R – La(1), Gd(2), Lu(3). (Reprinted with permission from [63] Copyright (2006) Pleiades Publishing Inc.)

Doping Influence on Fluorine-containing Phases

445

The percolation limit differs for various RE and depends on both the size factor and the cluster type existing in the structure [63]. In some solid solutions at high doping cation concentrations the conductivity starts to decrease and maxima may be observed on concentration dependences of conductivity [68–70]. Some examples are presented in Figures 14.21 and 14.22. From the standpoint of the defect region model, it means that, at high defect concentrations, some interstitial fluoride ions do not take part to the ionic transfer because they are incorporated in defect clusters. The charge transfer occurs only at the border of defect regions (at the interface cluster/fluorite matrix). Therefore at high defect region concentrations the conductivity falls. log σ [Ω–1cm–1]

Ea(eV)

0,5 –4

50 °C

0,4

–5 x 0

0,05

0,10

0,15

x 0

0,05

a

0,10

0,15

b

Figure 14.21 Concentration dependences of conductivity and activation energy at 323 K of Pb1xZrxF2þ2x solid solutions. (Reprinted with permission from [68] Copyright (1986) Elsevier Ltd.)

σ500/104(S/cm)

(a)

(b)

Ea (eV)

15

0.85

10

0.80

5

0.75

0

0.70 0

10 20 c (mol.% RF3)

30

0

10 20 c (mol.% RF3)

30

Figure 14.22 Concentration dependences of conductivity and activation energy at 500 K for BaF2-LnF3 solid solutions (Ln ¼ Gd -o-, Tb -D-, Y -&-) (Reprinted with permission from [69] Copyright (2008) Elsevier Ltd.)

446

Functionalized Inorganic Fluorides

It is possible to consider conductivity simultaneously in anion-excess and anion-deficit solid solutions since the anion sublattice defects – free vacancies, mobile interstitial ions and clusters – are always combined in such solid solutions. The clusters fix vacancies and interstitial anions in a rigid structure. For example, it was shown that concentration dependences of the conductivity and activation energy of Ba1tBitF2þt0.30O0.15 solid solutions (0.05 < x < 0.30) are monotonic and continuous for both anion-excess and anion-deficit composition regions (Figure 14.23) [71]. The conductivity maximum, observed in the anion-excess composition region, could be explained, similarly to Ba1tBitF2þt [72], by clustering defects. logσ473 K(Ω–1.cm–1)

Eσ(eV)

–2

1.4

–4

1.2

–6

1.0 Ba1–xBixF2+x

Ba1–xBixF2+x–0.30O0.15

–8

0.8

–10

0.6

X 0.1

0.2

0.3

0.4

Figure 14.23 Concentration dependences of conductivity and activation energy at 473 K for Ba1xBitF2þx0.30O0.15 and Ba1xBixF2þx solid solutions. (Reprinted with permission from [71] Copyright (2003) Elsevier Ltd.)

Conductivity maxima were also observed for anion-deficit solid solutions. For example, such a maximum was found out in Na0.10Nd0.90(O,F)2 solid solution for the composition Na0.10Nd0.90O0.96F0.88 ( ¼ 0.16) [73]. This fact is in full agreement with data obtained for some oxide fluorite-like solid solutions, for example ZrO2 – Yb2O3 [4]. The presence of defects clusters in disordered solid solutions was confirmed by direct experimental methods (neutron diffraction, 19F NMR, EXAFS, etc.) as well as indirectly, using results of X-ray and electron diffraction studies of ordered fluorite-like phases having close compositions [7]. It is clear that the cluster type, their concentration and transformations in nonstoichiometric phases influence their transport properties. But conductivity depends even more on the material’s microstructure – in particular its micro-heterogeneity.

Doping Influence on Fluorine-containing Phases

14.6.2

447

Nonstoichiometric Fluorites as Examples of Nanostructured Materials

It was suggested that the micro-heterogeneity appears in the disordered fluorite-like solid solution structure by means of ordering superclusters [43]. The superclusters accumulate R3þ dopant cations and form small areas. These ordered areas with structures similar to those of ordered phases are distributed in the matrix of non-distorted cubic fluorite. In Figure 14.24][43] two models of this micro-ordering in the disordered M2xRxF2þx phase (a,b) as well as an extended superstructure in ordered phases (c,d) are presented. In this figure the two-dimensional section of the cationic sublattice is shown. The single {M8[R6F69]} supercluster is presented in the left top corner of Figure 14.24a. Four R3þ cations can be seen in this section of the sublattice. Other cations of the supercluster, 8M2þ and 2R3þ, are located in neighbouring layers.

Figure 14.24 Models of M1xRxF2þx structure: a), c) - micro-ordering areas in nonstoichiometric crystals with x ¼ 0,15 and 0,40; b), d) ordered phases at x ¼ 0.1538 (hypothetic phase) and at x ¼ 0.4286 (real Ba4Yb3F17 phase [54]). Light circles denote R3þ, black ones, M2þ. (Reprinted with permission from [43] Copyright (2003) Pleiades Publishing Inc.)

Microordering areas are about 10–100 nm long. It is impossible to detect these randomly oriented nanometric areas by means of classical XRD because both the fluorite matrix and ordered inclusions have an almost identical cationic sublattice. Slight cation displacements

448

Functionalized Inorganic Fluorides

from the fluorite position (see the inset in Figure 14.24d) do not disturb the coherent concretion of various crystal lattice domains. But some modern experimental methods of structure investigation, for example high resolution electronic microscopy, enable microheterogeneity areas to be detected as ordered domains (similar to domains in oxygenconductive fluorites, see Figure 14.25 [74]). The suggested simplified model [43] of nonstoichiometric fluorite phases has some typical features: 1. The model deals with large isomorphous structure blocks of different cationic compositions but not with separate isomorphous ions. The maximum size difference of superclusters is 5.4 % (for fluorite matrix [Ba14F64] and RE superclusters [Ba8{Lu6F69}]). It leads to the removal of dimensional restrictions between M2þ and R3þ cations, which exist in the classic isomorphism model. 2. Separate superclusters group together and form micro-ordering areas (nanoparticles) retaining the fluorite cationic motive of the structure. The sizes of such areas are different ranging from several nanometers up to few microns.

Figure 14.25 High resolution image ({110}F) and selected area electron diffraction pattern ({110}F) recorded from Sm0.20Ce0.80O1.90 sintered body. Dashed line area means micro-domain with ordered structure. (Reprinted with permission from [74] Copyright (2002) Elsevier Ltd.)

The described model supposes that the micro-ordered areas can correspond even to hypothetical or metastable ordered phases, which are absent on the equilibrium phase diagram. For example, in the process of research of nonstoichiometric

Doping Influence on Fluorine-containing Phases

449

Na0.5xBi0.5þxO2x F2þ22x solid solutions by 19F NMR and electron diffraction methods, the presence of two basic cluster types (i.e. cuboctahedral (8:12:0) and triangular (1:0:3) [41,70,75]) has been shown. It is rather interesting information because there are no ordered phases with cuboctahedral clusters, similar to KBi3F10 [46], in the NaBiF4–BiOF–BiF3 [76] system. At the same time there exists an ordered phase with triangular clusters Bi(O,F)2,45 in this system [58]. The conductivity maximum in this cubic solid solution is observed for the structures with the maximum clusters variety and the highest unit cell parameters. Thus, according to [43], nonstoichiometric fluorite-like M2xRxF2þx phases with micro-ordering could be considered as a new class of nanostructured materials. These materials differ from known ones by the isostructurallity and coherency of the matrixes and nano-inclusions though they have different chemical composition. For a long time, micro-heterogeneity of these materials has not been detected because they had a single crystal outward appearance and the presence of nano-inclusions could not be detected by many diffraction methods. In addition the concentration dependence of their unit cell parameters was quite similar to the usual change of these parameters in solid solutions having wide range of homogeneity (up to 50 mol % RF3). The ionic conductivity is a characteristic of materials rather sensitive to changes of the defect structure and to micro-ordering. The model of micro-ordering developed for description of fluorite-like solid solutions and the defect region model used for description of many objects explain equally discrepancy between the observed charge carriers concentration and the quantity of anion defects [43]. Besides of this, defect area volumes (3000–4000 A3) estimated from the Ba2xRxF2þx conductivity data are very close to the associated supercluster size obtained by the micro-ordering areas model [43]. The fluorite-like nonstoichiometric phases models of micro-ordering (nano-inclusions) and the defect region model are also applicable to less commonly studied tysonite-like solid solutions and to nanocomposite materials having high ionic conductivity. For example, nanocrystalline BaF2:CaF2 composites prepared by high-energy ball milling have unexpectedly high dc conductivity of about 0.1 mScm1 at 450 K [77]. Evidently boundary effects have a crucial role not only for nanomaterials but for disperoids – eutectic compositions, fluorides with dielectric nano-phase dilution etc – as well [78–80]. There are not only maxima on concentration dependences of conductivity of nonstoichiometric phases but sometimes also minima in the area of very low RF3 concentrations (less than 2 %). For example, for Pb1xYbxF2þx the minimum was observed at x ¼ 0.005–0.02 at 293 K (depending on conditions of the sample thermal treatment) [10]. The appearance of these minima is usually connected to the introduction of structural defects with smaller mobility than the one of intrinsic defects into the crystal matrix.

14.7

14.7.1

Peculiarities of the Defect Structure of Nonstoichiometric Tysonite-like Phases Tysonite Structure, Tysonite Modifications and Anion Defects

The name of the tysonite structural type comes from the mineral tysonite, a solid solution of cerium and other RE fluorides in LaF3. There are two modifications of the tysonite

450

Functionalized Inorganic Fluorides

˚ and structure: a hexagonal phase I (P63/mmc, Z ¼ 2, cell parameters a ¼ 4.148 A ˚ c ¼ 7.354 A for Ce0.5La0.5F3 [81]) and a trigonal phase II (P3c1, Z ¼ 6, cell parameters ˚ , c ¼ 7.351 A ˚ for LaF3 [82]). a ¼ 7.185 A Hexagonal Modification. The structure I consists of layers of various composition that are perpendicular to the c axis (Figure 14.26) [81]).

c

F1 F2 La, Ce

Figure 14.26 View perpendicular to the c axis of the hexagonal modification of tysonite (I) for La0.5Ce0.5F3 [after 81]*

The first layer consists only of fluorine atoms (F1), which form a hexagonal grid similar to the carbon atom arrangement in graphite. These layers are situated at z ¼ 0.5 along the c axis. The second layer is a mixed one and consists of alternating Ln and F2 atoms (in Figure 14.26 lines connect neighbouring atoms to show the hexagonal motive of atoms in the layer). These layers are situated at z ¼ 0.25 and z ¼ 0.75 along the c axis. F1 and F2 atoms have different cationic environments. F1 is in the centre of a cation tetrahedron and F2 is surrounded by three cations and is situated in the centre of a triangle. The fluoride anion ratio is F1:F2 ¼ 2:1. The coordination number of the cation is 11. The irregular coordination polyhedron can be described as a trigonal prism with additional fluorine atoms at some distance from the centre of each face above its plane. In case of La0.5Ce0.5F3, La,Ce atoms have the following ˚ , three F2 fluorine atoms at 2.395 A ˚ and coordination environment: two F1 atoms at 2.353 A ˚ six more distant F1 atoms at 2,736 A. All these fluorine atoms form an irregular polyhedron. It was suggested that the hexagonal tysonite modification (I) becomes more stable when the anion packing density is reduced by the heterovalent replacement of RE atoms in LnF3 (Ln ¼ Gd–Er) by Mþ2 cations [83]. Trigonal modification. There are slight differences between both tysonite modifications. Translational symmetry of the cationic sublattice remains invariable but there are some changes in the anionic sublattice. In the trigonal modification (structure II) F2 position splits into two ones (F2 and F3) and some displacements of F1 atoms occur from their symmetric positions. A projection of the unit cell for the structure II onto the plane x0y is presented in Figure 14.27b and the corresponding part of the structure I projection onto the same plane is shown in Figure 14.27a.

* Hereinafter the positions F1 and F2 in the modification I are renamed (inversion of F1 and F2) for more convenient comparison of this structure with the tysonite modification II.

Doping Influence on Fluorine-containing Phases a

b

451

a

b

F1 F2 La, Ce a

F3 F2 F1 La b

Figure 14.27 Projection of the hexagonal (a) and trigonal (b) tysonite structures onto the x0y plane [after 81, 84, 85]

As a result Ln and F1 atoms take initial positions in the formed cell and F2 and F3 atoms are displaced from symmetry planes existing in the structure I. This displacement leads to the increase of the parameter of the unit cell of the modification II. The arrangement of atoms in the trigonal phase is also close to a layered structure but layers in this case are not flat. F2 atoms move out from the Ln–F plane, because of a very small F-F distance between fluorine and Ln-F layers, and take place not far from the new Ln-F3 plane. The ratio of atoms in these positions is F1:F2: F3 ¼ 12:4:2. Ln polyhedron is a strongly distorted trigonal prism with five additional fluoride ions ˚ (for LaF3) with nine over all faces. The Ln–F distances are within the limits of 2.42–3.00 A ˚ ˚ [85]. bonds having length of 2.42–2.64 A and two remaining ones of 3.00 A The following phases have the hexagonal tysonite structure I: the mineral tysonite, ThOF2 [86] and BiO0.1F2.8 [87] oxyfluorides and MxR1xF3x solid solutions (R ¼ Gd-Er; M ¼ Ca, Sr) [88]. Possibly, the high anion defects concentration (OF and VF) promotes stabilization of this modification. The limits of x in the MxR1xF3x solid solutions are 0.07 £ x £ 0.33 (phase B in Figure 14.17). The following phases adopt the trigonal tysonite structure II: La-Nd trifluorides, solid solutions on their basis MxR1xF32x (M ¼ Ca, Sr, R ¼ La - Nd) [88] and BaxBi1xF3x [44]. Phase transitions between tysonite modifications are quite possible. When heated, a transformation of phase II to more symmetric modification I has been detected for RF3 (R ¼ La-Eu) [89]. For SrxYb1xF3x [90] and Bi1xNdx(O,F)3 [22] solid solutions, a polymorphic transition was also found. It was shown for the BiO0,1F2,8–NdF3 system (Figure 14.28) that the phase transition temperature decreases with increasing the anion vacancy and oxygen ion concentrations in Bi1xNdx(O,F)3 solid solutions. In the tysonite-like Bi1xBax(O,F)3 [11] solid solution of both tysonite modifications exists (Figure 14.16). The anionic deficit () is 0.10–0.18 and 0.03–0.10 for modification I and II, respectively. p The unit cells parameters of these phases are connected by the relations: aI ¼ aII/ 3 and cI ¼ cII. The presence of superstructure reflections of the modification II was confirmed by electronic diffraction (ED) as shown in Figure 14.29.

452

Functionalized Inorganic Fluorides T,K X* 1280

phase I 923

I + II 873

phase II 823 BiO0.1F2.8

50%

NdF3

Figure 14.28 Part of the phase diagram for the BiO0,1F2,8–NdF3 system [22]. X* is the temperature of phase transition between tysonite modifications I and II for NdF3. (Reprinted with permission from [89] Copyright (2008) Pleiades Publishing Inc.)

Figure 14.29 ED image for the Ba0.03Bi0.97O0.04F2.89 phase (tysonite II) [11]

The tysonite structure is close to the structure of orthorhombic -YF3 [91] (Figure 14.30).* Orthorhombic BiF3 crystallizes also in -YF3 structural type [91]. BiF3 transfers under slight static pressure (15 – 3 kbar) to tysonite type structure [8]. * Rare-earth trifluorides crystallize in three structural types: the tysonite structure type, the hexagonal -YF3 structural type and orthorhombic -YF3 structural type [7].

Doping Influence on Fluorine-containing Phases

a)

b)

453

c)

Figure 14.30 Cationic coordination polyhedra: a) tysonite modification I, b) tysonite modification II and c) ß-YF3

Bi3þ cation (having a 6s2 lone electron pair) has a higher deformability than Nd3þ cation though both cations have close ionic radii. Therefore the polymorphic transition of BiF3 from -YF3 structural type to tysonite II structural type, similar to the transition in RF3 (R ¼ Sm-Tb) [7], and then to the tysonite I structural type, similar to the transition in RF3 (R ¼ La-Pm) [89], occurs under mild conditions. Similar phase transitions are observed also in presence of anion vacancies in BiF3 by heterovalent replacement of Bi3þ by M2þ or of F by O2. As a result anionic conductivity rises sharply. However at high concentrations of anion vacancies (and correspondingly, O0 F or M0 Bi defects), at almost invariable unit cell parameters, defect ordering occurs with formation of defect clusters and new phases. Micro-ordering and crystallization processes of the tysonite-like ordered phases reduce anionic conductivity of materials. Presented examples of interrelations of both tysonite modifications and phases with the ß-YF3-structural type are in good agreement with the (P,T)-diagram in Figure 14.31[7]. Thus, the joint influence of heterovalent doping by O2 ions and by a light-deformable cation causes structural changes similar to that occurring by temperature and pressure rise, up to formation of new ordered phases.

T

Temperature

α-UO3

ReO3

VF3

LaF3 tysonite

LaF3 cubic

β -YF3 Pressure

P

Figure 14.31 [7]. (P,T) schematic phase diagram of RE fluorides structural types [92] with an account of the data reported. (Reprinted with permission from [93] Copyright (2001) Bovis P. Sobolev.)

454

14.7.2

Functionalized Inorganic Fluorides

Ordered Tysonite-like Phases

In CaF2–RF3 (R ¼ Gd–Yb, Y) systems the formation of tysonite-type ordered phases was detected [94–96]. The phases were obtained by solid state synthesis from corresponding fluoride mixtures (1173–1373 K, 3 weeks) with quenching to room temperature following. All samples obtained contained two phases, an ordered tysonite-like phase and a nonstochiometric solid solution (hexagonal modification I). Based on the powder X-ray and electron diffraction data [96], the composition of the ordered phases was found as group is Cc or C2/c and the unit cell parameters of Ca3R7F27. The possible space ˚ , b ¼ 6.716 A ˚ , c ¼ 7.971 A ˚ , ¼ 119.48. Ca3Y7F27 are: a ¼ 19.803 A In the BiF3-Bi2O3 system (Figure 14.16) an ordered tysonite-like BiOyF32y (0.13 < y < 0.23) phase [34] was also found. On the X-ray powder patterns of this phase (III), splitting of hexagonal structure lines (tysonite I lines) was observed. The X-ray powder pattern of the phase III was indexed in monoclinic symmetry (space group C2/c, Z ¼ 4) by the homology method. The subcell parameters for the composition BiO0,19F2,62 are: ˚ , b ¼ 7.088(4) A ˚ , c ¼ 8.414(4) A ˚ and ¼ 119.80(4). a ¼ 4.112(1) A

14.8

Ionic Transfer in Tysonite-like Phases

At the present time tysonite-like fluorides belong to the group of the best fluoride-ion conductors. Their conductivities are higher than those of fluorite phases. The most striking example is the commercial solid electrolyte La0.992Eu0.008F2.992. The ionic transfer in tysonite phases is caused by the presence of anion vacancies in the structures. By heterovalent doping, the formation of additional anion vacancies occur and, accordingly, the conductivity is by one or two orders of magnitude greater than in the initial tysonite. This heterovalent doping could be carried out into the cation sublattice (for example, La1xEuxF3x), anion sublattice (for example, BiOyF3y) or simultaneously both sublattices (for example, Bi1xBax(O,F)3). The anion vacancy mechanism of conductivity for RE metal fluorides was confirmed by means of heterovalent doping of LaF3. A sharp increase of conductivity has been observed by doping LaF3 by MF and MF2 fluorides but its doping by ThF4 leads to a decrease of the conductivity [97]. The vacancy mechanism of conductivity was also confirmed by 19F NMR [98]. To determine the ionic transfer mechanism it is necessary to know which structure positions are occupied by anion vacancies. It was shown earlier (see 14.7.1) that there are two different modifications of tysonite. The main difference between them consists in the fact that the modification I contains two sublattices, F1 and F2, and the modification II contains three sublattices, F1, F2 and F3. For example, the 19F MAS NMR spectra of LaF3 and La1.99Sr0.01F2.99 demonstrate that the three crystallographically distinct fluoride ion sites have very different mobilities, and the activation energies, for the different fluoride ion jump pathways, increase in the order F1-1 < F1-3 < F1-2. [99] (Figure 14.32b). Studying the mechanism of charge transfer in tysonite-like structures researchers combine sometimes F2 and F3 positions together and consider only two fluoride

Doping Influence on Fluorine-containing Phases

455

sublattices [100,101]. The first sublattice is uniform and consists of F1 positions only, the second one is the joint F2,3 system with the following ratio of positions: F1:F2,3 ¼ 2:1. The results of 19F NMR studies of La1xSrxF3x (modification II) [102] show that the increase of x from 0 to 0.16 (i.e. the increase of Sr content and correspondingly of the anion vacancy concentration) leads to the decrease of the ratio of intensities of peaks corresponding to F1 and F2,3 from 2.00 (–0.05) to 1.80 (–0.05). Thus in La1xSrxF3x solid solutions, vacancies occupy mainly F1 positions. 14.8.1

Fluoride Ions’ Migration Paths in the LaF3 Structure

The 19F NMR studies of LaF3 (the modification II) [98,99,103–107] showed inhomogeneity of the dynamic behaviour of anion sublattice in tysonite-like compounds. It was found that F1 is a high-mobile subsystem and F2,3 is a slow-mobile one at low temperatures (Figure 14.32). F1 F2

24 C *

*

F3 *

*

102 C

F1

140 C

198 C

T (K) F2.3

400 360 340 320 300 290 280 270 260 250 220

237 C –19.7 ppm

266 C 10 kHz

100

a

50

0 ppm

–50

–100

b

Figure 14.32 19F NMR spectra of a LaF3 single crystal. (Reprinted with permission from [107] (a) Copyright (1999) Springer. 19F variable-temperature MAS NMR spectra of LaF3, collected with a spinning speed of 23 kHz (* ¼ spinning sideband). Reprinted with permission from [99] (b) Copyright (1997) American Chemical Society.)

Temperature dependence of the ionic conductivity in the LaF3 is presented in Figure 14.33 [108]. At low temperatures the ionic transfer occurs within the uniform fluorine sublattice F1 (range I). At higher temperatures, an exchange over all fluorine sublattices begins (range II). It results in an activation energy decrease from 0.34 eV (range I, along c-axis) to 0.18 eV (range II) at higher temperatures. At even higher temperatures the thermal intrinsic defect concentration in pure LaF3 becomes large enough and the range III is essentially governed by intrinsic mechanism. The activation energy in this range is 1.22 eV.

456

Functionalized Inorganic Fluorides 102

T

to c-axis (first run) II to c-axis (first run) to c-axis (third run)

T

σT (Ω–1 cm–1 K)

101 100 10–1 10–2 10–3

II

III

I

10–4 0.5

1.0

1.5

2.0 2.5 1000/T (K–1)

3.0

3.5

Figure 14.33 Temperature dependence of the ionic conductivity in the LaF3 single crystal for two orientations along c-axis and perpendicular to the c-axis. The Roman letters denote ranges of different conductivity mechanisms. (Reprinted with permission from [108] Copyright (2003) Elsevier Ltd.)

By the doping of LaF3 single crystals by SrF2 the ionic conductivity increases and the ranges I, II and III change their boundaries: the II, governed by extrinsic conductivity due to external doping becomes broader; the range III shifts to higher temperatures and practically disappears if the dopant concentration exceeds 1 % [109]. It was found in this paper the presence of five ranges responsible for different regimes of fluorine dynamics in LaF3-type fast ionic conductors. The combination of microscopic (NMR spectra, relaxation) and mesoscopic (diffusion, conductivity) data obtained in this research leads to consistent overall picture of fluoride ionic dynamics. From the temperature dependencies of the fluorine dynamics a qualitative picture of a hierarchy of thermally activated processes has been deduced. The aging phenomenon in pure and doped LaF3 was observed after several heatingcooling runs. This phenomenon becomes evident in gradual changing the temperature dependence of conductivity in heating-cooling cycles. The reason is the appearance of additional anion vacancies VF by introducing oxygen ions into fluoride samples due to pyrohydrolysis, interaction with the gas phase or with other surroundings [108,109]. Similar influence of heterovalent substitution both in anon and cation sublattices on the conductivity of tysonite-like solid Bi1xBax(O,F)3 solutions was described in [11]. According to Figure 14.33, a slight anisotropy of transport properties of LaF3 exists in range I (below 450 K); the values of conductivity measured along (||) and perpendicular (?) to the c axis are different. Same results were obtained by conductivity measurements of LaF3 single crystals [110,111]. A model describing the anion migration in tysonite-like compounds was suggested [111]. This model explains the anisotropy of transfer properties. According to this model there are four possible options for fluorine atom jumps denoted as p0, p1, p2 and p3 in Figure 14.34. p0 and p3 jumps correspond to the fluorine atom transfer between F1 positions. Namely, they are involved in the ionic transport at low temperatures. p1 and p2 jumps correspond to the fluorine atom transfer between F2,3 and F1 positions.

Doping Influence on Fluorine-containing Phases

457

The distances F1–F1 and F1–F2,3 are approximately equal but NMR studies showed that the fluoride ion lifetime in positions F2,3 is longer. It could be probably attributed to a lower free energy of these positions (higher coordinations with regard to La3þ) as well as to the lower number of possible jumps to nearest neighbour positions [111]. Thus, F-sites on B-planes act similarly to trap sites and F3 sites block the charge transport along c-direction.

A F1 F3

B La

F2

A p0

A

p3

B p2

p1

A

C

Figure 14.34 F- ions migration paths in LaF3. (Reprinted with permission from [111]. View of the tysonite structure perpendicular to the c-axis (?c); the different F sites are ordered in different layers (sequence AAB along the c-axis: A contains only F1, B only F2/F3 sites). Copyright (1997) Elsevier Ltd.)

Thus, at low temperatures, only p0 and p3 jumps occur. Moreover p3 jumps occur along the c axis and correspondingly do not contribute to ? conductivity. Both types of jumps take part in transfer along the c axis, explaining the transport property anisotropy in LaF3. The F transport along the c-axis is faster, because a short and direct pathway exists via F1 sites on the A-plane. Apart from this, a higher free energy of F ions on A-planes may support a relative increase of the F vacancy concentration on A-planes. This approach was recently developed by the authors of [109]. They succeeded in studies of superfast fluorine dynamics in LaF3 single crystals using several advanced NMR methods (spectroscopy, relaxometry, diffusiometry). From the high frequency relaxation data it was possible to gain short motional correlation times of down to almost 1011 s. The fluorine dynamics was described assuming equal jump probabilities among all neighbouring F1 sites. Three of them belong to the same F1 layer (plane perpendicular to c-axis) while the fourth site, displaced almost parallel to the c-axis, belongs to the next layer. The sum of the three intraplane jump vectors outbalances the interplane jump vector. It was concluded that fluorine dynamics within the F1 sublattice should be approximately isotropic. 14.8.2

Temperature Dependences of Ionic Conductivity and Anion Defect Positions

Similarly to LaF3, almost all tysonite-like phases (modification II) have an inflection in the temperature dependence of the conductivity (Figure 14.33, see the boundary between

458

Functionalized Inorganic Fluorides

ranges I and II). But this inflection is absent for the stoichiometric oxyfluoride ThOF2 [112]. The authors of [113], taking into account ordering anions in ThOF2, concluded that oxygen ions do not take part in the conductivity. This ordering means that all oxygen ions occupy positions in cation layers (F2,3 positions) and all F ions are between these layers (F1 positions). Another reason for this conclusion could be the low value of Ea (0.37 eV). Thus the ionic conductivity of ThOF2 is governed by the fluoride ion concentration and mobility. Accordingly, the transfer mechanism is identical at all temperatures (ionic transfer only through F1 positions and no inflection in the temperature dependence of the conductivity). There are no inflection in temperature dependences of the conductivity for R1xSrxF3x (R ¼ Tb, Ho, Y, Er, Tm, Yb, Lu; 0.04 < x < 0.33) having the hexagonal structure I. The Ea values of the conductivity of these nonstoichiometric phases with high concentration of anion vacancies are in the range 0.581(5) – 0.646(2) eV [84]. It was supposed that vacancies are mainly located in the F2 sublattice [83, 84, 114]. This assumption agrees with results obtained for the disordered tysonite-like Bi1xBax(O,F)3 solid solution [11] which exists in the form of both modifications I and II (Figure 14.16). It is assumed that positions F2 of the modification I are partially occupied by oxygen and vacancies [11]. As for the modification II, the vacancies are in F1 positions and oxygen ions are located in F2,3 positions. Probably, oxygen ions of the modification II do not take part in the conductivity and because they are located in the cations layers (positions F2,3) they hinder fluoride ion jumps along the c axis between two F1 positions (Figure 14.35) through cation-anion layers. VF1 La

2.63

F2,3 2.52

F1

2.51

VF1 C

Figure 14.35 example)

Conductivity mechanism in phases with the tysonite structure (LaF3 as an

Anion vacancies in cation-anion layers of the modification I reduce these hindrances. That is why a decrease of oxygen ion contents in modifications I and II has different impact on the conductivity: decrease of the conductivity of modification II and increase of the conductivity of I. It could be seen for example by observing the concentration dependence of conductivity of Bi0,96Ba0,4OyF2,962y samples in Figure 14.5 (y < 0.7, modification II, y > 0.7, modification I).

Doping Influence on Fluorine-containing Phases

459

Because all anion positions of both Bi1xBax(O,F)3 modifications [34] are occupied by mobile F ions (contrary to the stoichiometric oxyfluoride ThOF2), inflections in temperature dependences of the conductivity at 400–430 K appear (Figure 14.36). The reason for this phenomenon consists in the involvement of anion positions of the F2,3 sublattice in transport processes. Both regions of this dependence are well described in terms of the Arrhenius–Frenkel equation: T ¼ 0 exp[Ea/kT]. At low temperatures, Ea1 0.56–0.63 eV and above Tc, the ions of both sublattices are mobile and the conductivity activation energy falls down to Ea2 0.12–0.16 eV.

1

Bi0,96 Ba0,04 Oy F2,96-2y

lgσT S CM–1 K 0 y = 0,15 –1

y = 0,0375 y = 0,0675

–2 y = 0,0875 –3 2,0

2,4

2,8

3,2

3,6

1000/T,K

Figure 14.36

Temperature dependencies of conductivity of Bi0,96Ba0,4OyF2,962y samples [34]

The conductivity of tysonites, as well as that of fluorites, is affected by the presence of polarizable cations with a lone electron pair. The temperature dependences of conductivity of modifications I and II, for Bi1xNdx(O,F)3 solid solutions (Figure 14.28) are presented in Figure 14.37[22]. The conductivity of solid solutions under investigation increases, in both phases, when the bismuth ions content (and correspondingly, deficit of anions) increases. But the conductivity is higher by one or two orders of magnitude in modification I which contains more anion vacancies.

14.8.3

Concentration Dependences of Ionic Conductivity in Tysonite-like Solid Solutions

Similarly to fluorite-like solid solutions, maxima on the concentration dependence curve of conductivity were found for heterovalent doping tysonite-like solid solutions with I and II structures (Figure 14.38 [115] and 14.39 [34]).

460

Functionalized Inorganic Fluorides

1.6 1

1.8

2.0

2.2

1000/T, K–1 2.4 2.6 2.8

3.0

3.2

3.4

Bi1–x Ndx (O, F)3–d

0

x = 0.3 x = 0.9

–1

x = 0.2 x = 0.8

x = 0.1 x = 0.7

–2 –3 –4 –5 –6 logσ [S/cm]

Figure 14.37 Temperature dependence of conductivity for Bi1xNdx(O,F)3d tysonite solid solutions (modification I for 0 < x < 0,3 and modification II for 0.7 < x < 1). (Reprinted with permission from [22] Copyright (2008) Pleiades Publishing Inc.)

2

lgσ [S/m]

4

0

–2

3

–4 Ea, eV 0.6

2 1

0.4

0.2 0

5 10 SrF2, mol.%

15

Figure 14.38 Concentration dependence of conductivity and activation enthalpy for La1xSrxF3x. 1,3: T < Tc. 2,4: T > Tc. (Reprinted with permission from [115] Copyright (1994) Kristallographia Publishing Company.)

Doping Influence on Fluorine-containing Phases

461

It was assumed that, as with fluorites, the transfer model, taking into account clustering processes, percolations and charge carrier interactions, is suitable for La1xSrxF3x solid solutions (modification II) [115]. With replacement of La3þ by Sr2þ extrinsic fluorine vacancies appear in the F1 sublattice. By doping, the concentration of Sr2þ ions increases and these ions seem to aggregate with anion vacancies forming clusters similar to associates of defects in fluorite-like nonstoichiometric phases M1xRxF2þx. The study of clustering in disordered solid solutions with the tysonite structure is just beginning. But the ordering of tysonite-like Ca3R7F27 [96] and III-BiOyF32y [34] phases has been already described. It could be supposed that, at least a part of extrinsic fluorine vacancies, in disordered phases, is located in structurally distorted areas of the anion sublattice around such clusters. It results in the mobility increase of charge carriers and then in an increase of the conductivity of nonstoichiometric solid solutions in comparison with stoichiometric MF3. When the doping concentration increases, the number of defect regions increases and, at a concentration called the percolation limit, xper, the defect regions combine themselves into the joint conductivity channel. The following clustering and probably local cluster ordering (similar to fluorites, see Section 14.5.2) in the structure leads to a decrease of the charge carrier concentration and to a blocking of a part of the conductivity channels. It negatively affects the transport properties of these phases. All these factors lead to the appearance of a maximum on the curve of conductivity, ¼ f(x) at x ¼ 0.05 (Figure 14.38). Similar extremes were observed on concentration dependences of transfer characteristics of tysonite-like MxR1xF32x (M ¼ Ca, Sr, R ¼ La-Lu) solid solutions [115].

Ea, eV

Bi1-x Bax O0,07 F2,86-x

16 σ*10 S CM–1

Bi1-x Bax O0,07 F2,86-x

0,64

T = 360 K

–6

12 0,56 8 0,48

4

0,40

0 0,00

0,02

0,04

0,06

X, mol. %

0,08

0,10

0,00

0,02

0,04

0,06

0,08

0,10

X, mol. %

Figure 14.39 Concentration dependences of conductivity and activation energy (T < Tc) for the solid solution Bi1xBaxO0.07F2.86x. (the homogeneity region is between the dotted line) [34]

A maximum on concentration dependences of the conductivity at t 0,07 was observed for the tysonite-like Bi1xBax(O,F)3 solid solution with the structure I (Figure 14.39). The discovery of the ordered tysonite-like BiOyF32y phase (0.13 < y < 0.23) [34], with close composition, could serve as an indirect confirmation of the clustering mechanism described

462

Functionalized Inorganic Fluorides

above. The formation of pair defects VF•-BaBi0 , VF•-OF0 could be assumed as the simplest clusters. Information about large-scale cluster formation in tysonite-like solid solutions is absent because there are no structural data about ordered tysonite-like phases.

14.9

Conclusions

The influence of different doping types on the defect structure and conductivity of fluorine-containing phases with fluorite and tysonite structures has been discussed in the current review. The ‘defect-region’ model including clustering and percolation phenomena has been used for describing the ionic transfer features in solid solutions. The heterovalent cation doping in fluorites (and probably also in tysonites) leads to the formation of micro-disordered solid solutions with wide homogeneity regions which could be considered as nano-structured materials. The radius ratio of dopant and matrix cations does not play a very important role in process of such doping because of the block isomorphism. Phase diagrams and structures of ordering phases are of great importance for interpretation of temperature and concentration dependences of transport properties as well as for the structure of nonstoichiomentric solid solutions. The doping by cations with an active lone electron pair or the use of two of such cations (of matrix and dopant) leads to an increase of the ionic conductivity. It is probably due to crystal structure and conductivity channel distortions. The heterovalent anion doping of fluoride phases by oxygen ions or simultaneous heterovalent replacement of both anions and cations is an effective method for additional influence on the defect structure and transport properties. The joint effect of easily deformable cations with a lone electron pair (BiF3 matrix) and heterovalent doping of fluorides by oxygen ions is similar to the effect of an increase in temperature and pressure. It has been shown that pyrohydrolysis proceeds spontaneously in fluorides and could lead to crucial changes of their properties or to appearance of new ordered oxyfluoride phases. Nanostructured fluorides and oxyfluorides, described in this review, could be considered as prospective solid electrolytes. Other promising substances for these applications are super-fine (eutectic) mixtures of fluorine-conductive phases and composites consisting of fluorine-conductive phases with inert nanoimpurities. The influence of doping on the defect structure and transport properties of fluorinecontaining phases could be useful for targeted searching for new solid electrolytes.

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[86] A. Rhandour, J.-M. Re´au, S. Matar, P. Hagenmuller, Influence de l’oxyge`ne sur les proprie´te´s cristallographiques et e´lectriques des materiaux de type tysonite, J. Solid State Chem., 64, 206–216 (1986). [87] A. Morell, B. Tanguy, J. Portier, Le syste`me Bi2O3–BiF3, Bull. Soc. Chim. Fr., 7, 2502–2504 (1971). [88] B. P. Sobolev, V. B. Alexandrov, P. P. Fedorov, K. B. Seiranian, N. L. Tkashenko, Variablecomposition phases with LaF3–type structure in the MF2-(Y,Ln)F3 system, Part IV: X-ray characteristics, heterovalent isomorphic substitutions, Kristallografiya, 21, 96–105 (1976) (in Russian). [89] O. Greis, M. S. R. Cader, Polymorphism of high-purity rare earth trifluorides, Thermochim. Acta, 87, 145–150 (1985). [90] S. N. Achary, A. K. Tyagi, J. Ko¨hler, Stabilization of high temperature tysonite type compounds in the Sr-Yb-F system-synthesis and Rietveld refinement, Mater. Res. Bul., 36, 1109–1115 (2001). [91] A. K. Cheetham, N. Norman, The Structures of Yttrium and Bismuth Trifluorides by Neutron Diffraction, Acta Chem. Scand., A 28, 55–60 (1974). [92] N. A. Bendeliani, Phase transformations of trifluorides of transition metals under high pressure, Izvestiya AN SSSR, Neorgan. Mater., 20, 1726–1729 (1984) (in Russian) [93] J. Ravez, A. Mogus- Milankovic, The ReO3-related structure ferroelastic fluorides, Jpn. J. Appl. Phys., 24, Suppl. 24–2, 687–689 (1985). [94] B. P. Sobolev, E. G. Ippolitov, About phase composition systems CaF2–YF3, SrF2–YF3, BaF2-YF3, Neorg. Mater., 1, 362–368 (1965) (in Russian). [95] L. S. Garashina, E. G. Ippolitov, B. M. Zhigarnovskii, B. P. Sobolev, Phase composition systems CaF2–YF3, in Investigation of Natural and Artificial Processes of Mineral Formation, Moscow, Nauka, 289–294 (1966) (in Russian). [96] D. J. M. Bevan, O. Greis, Fluorine-deficient tysonite-type solid solutions (Ca,Ln)F3 and related superstructure phases Ca3Ln7F27, Rev. Chim. Miner., 15, 346–359 (1978). [97] A. Roos, D. R. Franceschetti, J. Schoonman, The small-signal ac response of La1xBaxF3x solid solutions, Solid State Ionics, 12, 485–491 (1984). [98] M. Izosimova, A. I. Livshits, V. M. Buznik, P. P. Fedorov, E. A. Krivandina, B. P. Sobolev, Diffusion mechanism of F-ions in solid electrolytes with tysonite structure, Fiz. Tverd. Tela (S.-Petersburg), 28, 2644–2647 (1986) (in Russian). [99] F. Wang, C. P. Grey, Probing the Mechanism of Fluoride-Ion Conduction in LaF3 and StrontiumDoped LaF3 with High-Resolution 19F MAS NMR, Chem. Mater., 9, 1068–1070 (1997) [100] A. Roos, F. C. M. van de Pol, R. Keim, J. Schoonman, Ionic conductivity in tysonite-type solid solutions La1xBaxF3x, Solid State Ionics, 13, 191–203 (1984). [101] A. I. Livshits, V. M. Buznik, P. P. Fedorov, B. P. Sobolev, NMR Study of Anion Mobility in Defect Phases of the Fluorite and Tysonite Structure in the CaF2–LaF3 System, Neorg. Mater., 18, 135–139 (1982) (in Russian). [102] A. F. Privalov, H-M. Vieth, I. V. Murin, Nuclear magnetic resonance study of superionic conductors with tysonite structure, J. Phys.: Condens. Matter, 6, 8237–8243 (1994). [103] K. Lee, A. Sher, F19 Nuclear Magnetic Resonance Line Narrowing in LaF3 at 300 K, Phys. Rev. Lett., 14, 1027–1029 (1965). [104] M. Goldman, L. Shen, Spin-Spin Relaxation in LaF3, Phys. Rev., 144, 321–331 (1966). [105] A. G. Lundin, S. P. Gabuda, A. I. Livshits, Nuclear Magnetic Resonance and superfine interaction in crystals with tysonite structure, Fiz. Tverd. Tela (S.-Petersburg), 9, 707–710 (1967) (in Russian). [106] N. I. Sorokin, M. V. Fominykh, E. A. Krivandina, Z. I. Zhmurova, B. P. Sobolev, Ion transport in R1xSrxF3x (R ¼ La-Yb, Y) solid solutions with a LaF3 (Tysonite) structure, Crystallography Reports, 41, 292–301 (1996). [107] A. F. Privalov, I. V. Murin, Ion-motion disorder in a tysonite superionic conductor from 19F NMR data, Physics of the Solid State, 41, 1482–1485 (1999). [108] V. V. Sinitsyn, O. Lips, A. F. Privalov, F. Fujara, I. V. Murin, Transport properties of LaF3 fast ionic conductor studied by field gradient NMR and impedance spectroscopy, J. Phys. Chem. Solids, 64, 1201–1205 (2003).

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[109] F. Fujara, D. Kruk, O. Lips, A.F. Privalov, V. Sinitsyn, H. Stork, Fluorine dynamics in LaF3-type fast ionic conductors – Combined results of NMR and conductivity techniques, Solid State Ionics, 179, 2350–2357 (2008). [110] A. Roos, A. F. Aalders, J. Schoonman, A. F. M. Arts, H. W. de Wijn, Electrical conduction and 19 F NMR of solid solutions La1xBaxF3x, Solid State Ionics, 9, 571–574 (1983). [111] C. Hoff, H.-D. Wiemho¨ffer, O. Glumov, I. V. Murin, Orientation dependence of the ionic conductivity in single crystals of lanthanum and cerium trifluoride, Solid State Ionics, 101–103, 445–449 (1997). [112] P. Laborde, J.-M. Re´au, Mise en e´vidence et e´tude des proprie´te´s de transport de nouvelles solutions solides oxyfluore´es appartenant au syste`me ternaire BiF3-BiOF-ThO2, J. Solid State Chem., 72, 225–233 (1988). [113] A. Rhandour, J.-M. Re´au, S. Matar, P. Hagenmuller, Influence de l’oxyge`ne sur les proprie´te´s cristallographiques et e´lectriques des materiaux de type tysonite, J. Solid State Chem., 64, 206–216 (1986). [114] N. I. Sorokin, E. A. Krivandina, Z. I. Zhmurova, B. P. Sobolev, M. V. Fominykh, V. V. Fistul’, Superionic conductivity of the heterovalent solid solutions R1xMxF3x (R ¼ REE, M ¼ Ca, Ba) with tysonite-type structure, Physics of the Solid State, 41, 573–575 (1999). [115] N. I. Sorokin, B. P. Sobolev, Conductivity of single crystal solid solutions La1xSrxF3x (0 £ x £ 0.15) with tysonite-type structure, Kristallografiya, 39, 889–893 (1994) (in Russian).

15 Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups Yoshiaki Matsuo Department of Materials Science and Chemistry, Graduate School of Engineering, University of Hyogo, 2167 Shosha Himeji, Hyogo, 671-2201, Japan

15.1

Introduction

Photophysical properties and photochemical reactions of photoactive organic molecules in solid media may differ from those in homogenous solutions, in terms of the life time of the excited state, efficiency of radiationless quenching, diffusion of excited molecules, etc. [1, 2]. Accordingly, the quantum yield of fluorescence, and the distribution and stereochemistry of the photochemical products can be greatly changed. Therefore, the study of the photoprocesses of organic molecules in solid media is of interest because it could yield various applications such as solid dye laser, nonlinear optics, reaction media for controlled photochemical reactions and so on. Various solid materials have been used as hosts of photoactive organic molecules such as zeolties, mesoporous silicas, layered materials, etc. The advantage of layered materials to use as the host materials of photoactive organic molecules is that they possess twodimensional and expandable interlayer space, which can accommodate a wide variety of

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Ó 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

470

Functionalized Inorganic Fluorides

organic species. This type of reaction is called intercalation. The orientation and aggregation of the intercalated organic molecules can be controlled by designing and selecting both the hosts and guests, or adding third components. In addition, layered materials can be exfoliated to individual layers in appropriate solvents, giving so-called nanosheet solutions. Integration of these nanosheets can provide heterogeneous structure, thin films, disordered or highly ordered orientation of layers, etc, which may enhance the function of layers. On the other hand, perfluoroalkyl groups are known to show various interesting properties such as thermal, adhesive, frictional, electrical and photophysical properties, which mainly originate from the high electronegativity of the fluorine atom and the unique nature of C-F bonding [3]. Table 15.1 shows some physical properties of C-F bonding in comparison with those of C-H bonding [4–8]. The polarizability of C-F bonding is similar to that of C-H bonding, while the weight is much larger. This leads to the lower refractive index and higher transparency of the materials containing perfluoroalky groups [9, 10]. The materials containing perfluoroalkyl groups are, therefore, used as optical fibres, display coatings, etc. The perfluoroalkyl environment is also interesting from the viewpoint of microenvironments for chemical reactions because weak solventsolute interactions and relatively higher solute-solute interactions are expected [11, 12]. The micropolarity of the perfluoroalkyl environment can be estimated from the intensity ratio of the fluorescence peaks at 373 and 384 nm of pyrene (I1/I3 ratio) and is lower than that of the alkyl environment, as revealed by the smaller ratio of 0.50 in perfluoromethylcyclohexane than that (0.56) in methylcyclohenxane [13]. The other important feature of perfluoroalkyl groups is their high stability. The bond energy of C-F bonding is very large and the larger fluorine atoms hinder the carbon atoms, avoiding attack from other chemicals.

Table 15.1

Physical properties of C-F bonding in comparison with those of C-H bonding

Bond length/pm * Bond energy/kJmol1** Polarizability/1024 cm3 Vibration energy (stretching)/cm1 Volume of CX3 (X ¼ H or F)/106 pm3*** Volume of CX2 (X ¼ H or F)/106 pm3***

C-H

C-F

Ref

106–110 439 0.67 2850–3060 13.7 10.3

132–143 481 0.68 1000–1400 21.3 15.3

4 5 6 7 3, 8 3, 8

*: H or F atoms are bonded to sp3 carbons. **: H or F atoms are bonded to methyl groups.

The above facts strongly indicate that the layered materials hybridized with perfluoroalkyl groups would provide the environment with specific functions when photo-active organic molecules are included in them. In this chapter, synthesis of hybrid layered materials containing perfluoroalkyl groups and the incorporation of organic molecules into the resulting materials as well as the photophysical and photochemical properties of the introduced organic molecules are summarized.

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups

15.2

471

Preparation and Properties of Intercalation Compounds Containing Perfluoroalkyl Groups

15.2.1

Preparation

Four types of synthetic methods are known in order to introduce perfluoroalkyl groups into layered materials as shown in Figure 15.1, including (I) ion exchange reaction using ions containing perfluoroalkyl groups, (II) acid/base reaction with amines possessing perfluoroalkyl groups, (III) silylation by silylating reagents with perfluoroalkyl chains and (IV) oxidative intercalation of anions containing them. The layered materials that have been used as the hosts of perfluoroalkyl groups, the methods employed for intercalation and the compositions of the resulting intercalation compounds are summarized in Table 15.2 Layered materials such as clay minerals possess ion exchange capability and the ions in their interlayer space can be exchanged by foreign ions simply by immersing in an appropriate solution containing them as shown in Figure 15.1 (I) [14]. The exchangeable ions in clay minerals are usually metal cations, therefore, the amount of them per gram is

– – H+

–

H+

H+

–

H+

–

+

+

NH3

(II) Acid/base

–

Rf NH3

Rf

Rf

+

Rf

+

NH3

–

NH2

NH3

–

Ion exchange – M+

–

M+

M+

–

(I) Ion exchange

M+

Rf Rf

– + R

–

–

+

+ R

R

–

– + Rf

Rf

+

+

–

–

Rf

Ion exchange

+

+

– +

R

–

Rf

Rf

+

Rf

Si

Si

(III) Silylation

Si

Rf

Rf

Si

Si Cl

Figure 15.1 Schematic illustration of the introduction of perfluoroalkyl groups into layered materials. (I): ion exchange reaction, (II): acid/base reaction, (III): silylation and (IV): oxidative intercalation

472

Functionalized Inorganic Fluorides (IV) Oxidation

+

(Chemically or Electrochemically)

+

+

+

+ +

– Rf

Rf

+

+

+

–

–

– Rf

Rf Rf

Rf

Rf

–

–

–

+

+

+

Rf R

:

Perfluoroalkyl chain

:

Alkyl chain

M+

:

Metal ion

: A layer of host materials

Figure 15.1 (Continued)

called cation exchange capacity (abbreviated as CEC). The degree of ion exchange can be controlled by the concentration of cations in the solution. The amount of introduced perfluoroalkyl groups sometimes exceeds the CEC because neutral salt of surfactant cations combined with the counter anion can be also included between the layers by hydrophobic interaction as shown in Figure 15.2. This phenomenon is called intersalation and this was observed for surfactant-intercalated saponite clay shown in Table 15.2 [15]. On the other hand, in case of anion exchangeable layered materials, such as layered double hydroxide (abbreviated as LDH), anions containing perfluoroalkyl groups can be introduced between the layers of them. This type of reaction has been used for the removal of perfluoro-octane sulfonate or carbonate from water [16]. In order to intercalate large ions, a two-step ion-exchange method is performed. In case of the intercalation of large trans-[2(2,2,3,3,4,4,4-heprtafluorobutylamino)ethyl]-{2-[4-(4hexylphenylazo)-phenoxy]-ethyl}dimethylammonium ion (hereafter C3F-azoþ) ions into K4Nb6O17, potassium ions were first exchanged by methyl viologen ions and then, C3F-azoþ ions were reacted [17]. The increase of the interlayer spacing during the first ion exchange process facilitates the access of larger cations to the interlayer space in the second ion exchange process. The second method is the acid/base reaction used for the layered materials with high layer charges in which interlayer cations are tightly bound to the layer and, accordingly ion exchange reaction is difficult. The examples of this type of hosts are transition metal

Table 15.2

Preparation and compositions of intercalation compounds containing perfluoroalkyl groups

Host

Reagents

Saponite*

CnF2nþ1CONH(CH2)2Nþ(CH3)2(C16H33)Br (abbreviated as Ion CnF-S (n: number of carbons in perfluoroalky chain, 1–3)) exchange C8F17SO3H Ion exchange C8F17COOH

Layered double hydroxide (Abbreviated as LDH) K4Nb6O17 Graphite oxide (C8O4.0 H3.0****; abbreviated as GO) Magadiite (Na2Si14O29 nH2O) Graphite

C3F-AzoþBr*** C7F15CH2NH2 C6F13C2H4SiCl3

Method

Composition <4.4 CEC **

10

–

11

–

11

Ion – 12 exchange**** Acid/base (C7F15CH2NH2)0.38–0.62GO 19 Silylation (C6F13C2H4OH)0.45GO(C4H9NH2)0.40 27

C6F13C2H4Si(CH3)2Cl C6F13C2H4SiCl3

Silylation

CnF2nþ1SO3H (n ¼ 4, 6, 8)

Oxidative CxC4F9SO3d(C4F9SO3H) x ¼ 24,28, intercalation d¼2.5, 2.8 – Oxidative – intercalation CxC10F21SO3 x ¼ 20 CxC2F5OC2F4SO3 x ¼ 22 CxC2F5(C6F10)SO3 x ¼ 25 CxC8F17SO3dF x ¼ 9.6–136, d~4

LiN(SO2C2F5)2 /K2MnF6 LiN(SO2CF3)(SO2C4F9)/K2MnF6 KC10F21SO3 /K2MnF6 C2F5OC2F4SO3H/K2MnF6 KC2F5(C6F10)SO3 LiC8F17SO3

Ref

[(C6F13C2H4Si(CH3)2)1.6H0.4]Si14O29 [(C6F13C2H4SiOH)2]2.3Si14O29

* [(Si7.20Al0.80)(Mg5.97Al0.03)O20(OH)4]0.77(Na0.49Mg0.14)þ0.77 ** CEC: Cation exchange capacity; 0.997 meq/g. *** trans-[2(2,2,3,3,4,4,4-heprtafluorobutylamino)ethyl]-{2-[4-(4-hexylphenylazo)-phenoxy]-ethyl}dimethylammonium bromide. **** The potassium ions in K4Nb6O17 were first exchange by methyl viologen cations and then, they were exchanged by C3F-Azoþ. ***** The composition of graphite oxide greatly varies depending on the synthetic methods and ambient humidity.

26 58 33–35 37 37 38 38 38 39–41

474

Functionalized Inorganic Fluorides

– + Rf

+ –

Figure 15.2 compounds

– + Rf

Rf

+ –

Rf

X–

Rf

– +

+ Rf

Rf

X–

– +

+ Rf f

+ –

Rf – R X + + –

Rf

Schematic drawing for the intersalation of neutral ion pairs into intercalation

oxo-acids, such as H2Ti4O9 [18], HTiNbO5 [19], H0.7Ti1.83O4 [20], H0.5MoO3 [21], HNbWO6 [22] and HNbMoO6 [23]. In this case, the interlayer ions are first replaced by proton by acid treatment and then bases such as amines with perfluoroalkyl groups can be reacted. Perfluoro-octylamine was successfully intercalated into graphite oxide (abbreviated as GO) [24], which possesses a large number of acidic hydroxyl groups [25] and most of the intercalated amines are protonated to form ammonium ions in it [26]. The third method is the silylation of layered materials [27–35] using silylating reagents containing perfluoroalkyl groups, such as perfluoroalkylchlorosilanes, in organic solvents as shown in Figure 15.1 (III). In this case, the layered materials are first hydrophobized by replacing the interlayer cations or protons with hydrophobic species such as cationic surfactants or alkylamines and then the resulting materials are reacted with silylating reagents such as perfluoroalkylchlorosilanes in organic solvents. The hydrophobized layered materials are exfoliated in organic solvents, which facilitates the access of silylating reagents to the reaction sites. The perfluoroalkyl groups are covalently attached to the layers via Si-O bonding. The addition of a base to the reaction system is sometimes needed in order to scavenge the hydrogen chloride molecules that are formed as byproducts and destroy the resulting silylated products [36, 37]. The last method is oxidative intercalation and this method is mainly used for the intercalation of anions containing perfluoroalkyl groups into graphite [38–46]. Graphite is oxidized electrochemically in appropriate electrolyte solutions or by using oxidizing reagents such as K2MnF6 and anions with perfluoroalkyl groups are intercalated into oxidized graphite in order to compensate the charge as shown in Figure 15.1(IV). Compositions and stage structure of the resulting intercalation compounds are controlled by the applied potential or oxidizing reagents. Graphite intercalation compounds show various unique physical and chemical properties, however, are not usually optically transparent in the region of visible light, therefore, has not been used for the host of photo-functional molecules so far. However, recently, it has been reported that thin graphene films with thickness of 10 nm are highly transparent and conductive [47] and this would be interesting for the host of perfluoroalkyl groups. Not many reports on the preparation of intercalation compounds containing perfluoroalky groups have been provided. However, the other layered materials that are not shown in Table 15.2 and possess ion exchange capability or acidic hydroxyl groups are also expected to react with the reagents containing perfluoroalkyl groups in similar ways shown above.

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups

15.2.2

475

Exfoliation and Film Preparation

The layered materials are exfoliated into individual layers in appropriate solvents, forming so called ‘nanosheet’ solution. In aqueous systems, various inorganic layered materials such as clay minerals, titanates, manganates, niobates, etc., are known to exfoliate [48–57]. On the other hand, layered materials hydrophobized by organic species such as alky or perfluyoroalkyl chains are exfoliated in organic solvents. The exfoliation of hydrophobized layered materials is usually judged by observing whether they give a transparent solution in organic solvents with more than two different refractive indexes [33]. Transparent dispersions of [(C6F13C2H4Si(OH)2)]2.3Si14O29 were obtained in both ethylacetate and chloroform with different refractive indexes of 1.37 and 1.47, indicating the exfoliation of [(C6F13C2H4Si(OH)2)]2.3Si14O29 [58]. On the other hand, silylated magadiite without perfluoroalky groups dispersed in less polar solvents such as toluene. The stability of the exfoliated layers in aqueous solution can be described well by the Derjaguin-Landau-Verwey-Overbeek (DLVO) theory [59, 60], however, this theory cannot be applied for layered materials containing perfluoroalky groups without charge in organic solvents. The steric stabilization theory [61–64] which was originally developed to describe the stabilization of particles by polymers in nonpolar solvents would be more suitable for this system. In the present case, for the stability of the dispersion of hydrophobized layered materials, not only the hydrophobic interaction between solvent and alky or perfluoroalkyl chains but also the interaction between solvent and polar surface of layers seems important. Exfoliation of [(C6F13C2H4Si(CH3)2)1.6H0.4]Si14O29 in relatively polar 2-(perfluorhexyl)ethanol is also reported [26]. Further investigation is needed in order to understand the stability of the dispersion of hydrophobized layered materials. It is possible to obtain transparent thin films from the nanosheet solution by the cast method. An appropriate amount of nanosheet solution is put on the substrate and then it is simply allowed to stand until the solvent evaporates. Figure 15.3 shows the photograph and SEM images of the surface of the thin film of [(C6F13C2H4Si(OH)2)]2.3Si14O29 prepared from the nanosheet solution of the mixture of chloroform/ethylacetate [58]. The surface of the film became smoother when n-hexadecylamine was intercalated into [(C6F13C2H4Si(OH)2)]2.3Si14O29 and solvent was evaporated slowly. Starting from the nanosheets of intercalation compounds containing perfluoroalkyl groups, the other methods such as Langmuir-Blogdett technique [65, 66], layer-by-layer deposition [67, 68] which have been used for the film preparation of surfactant-intercalated clay, would be also applicable. B

C

50 μm

2 μm

Figure 15.3 Photograph and SEM images of the surface of thin film of magadiite silylated by 2-[hexyl]trichlorosilane. Adapted from Ref. 58

476

15.2.3

Functionalized Inorganic Fluorides

Introduction of Photofunctional Molecules

Three types of methods have been reported for the introduction of photofunctional molecules into layered materials containing perfluoroalkyl groups. The first is the ion exchange method, using the ions with both perfluoroalkyl groups and photo-functional moiety such as C3F-azoþ shown in Table 15.2. In this case the photofunctional molecules are bonded to the layered materials by ionic bonding. The second one is the exfoliationrestacking method. When neutral dye molecules are added to the nanosheet solution of layered materials containing perfluoroalyl groups in organic solvents and the resulting dispersion is cast on appropriate substrates, the exfoliated layers restack and the dye molecules are included in the obtained film by hydrophobic interaction. When starting from the aqueous dispersion of clay, by adding dye molecules in this dispersion, ion exchange reaction, restacking of exfoliated layers and inclusion of dye molecules occur at one time. The last one is the covalent attachment of dye molecules to silylated layered materials containing perfluoroalkyl groups. The procedure is schematically illustrated in Figure 15.4. An example of the above procedure is the covalent attachment of pyrene onto the layers of [(C6F13C2H4Si(OH)2]2.3Si14O29 shown in Table 15.2. This material contains silanol groups originated from the hydrolysis of Si-Cl groups, which were not used for silylation of magadiite layers. These silanol groups in this material were first silylated by 3-aminopropyltriethoxysilane, resulting in the introduction of amino groups. Then, an active ester containing pyrene shown in Figure 15.5 (1-pyrenebutanoic acid succimidyl ester), which selectively reacts with amino groups was added, resulting in the covalent attachment of pyrene chromophores by amide bonding [58]. Physically adsorbed dyes onto

F3C CF2

F2C H2N F3C

H2N

CF2 F 2C CF2 F2C CF2

C 2 H5 O

C2H5O Si OC2H5

F2C

O HN

CF2 O O

Si O

O

C2H5O C2H5O

Si

F3C CF2 F2 C CF2 F2 C CF2 O Si OH O

In N,N-dimethylformamide Silylated layered materials containing amino groups

Si OH

O N

Si OH O

C2H5O

HO O

C2H5O

CF2

Pyrene-attached silylated layered material

Toluene, reflux

Silylated layered materials F3C CF2

F2C

N C2H5O Si OC2H5 OC2H5

CF2 N F2 C CF2 C2H5O Si C2H5O O Si OH O

Pyrene-attached silylated layered materials

Figure 15.4 Schematic drawing for the procedures of covalent attachment of pyrene to silylated magadiite containing perfluoalkyl groups. Adapted from Ref. 58

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups

477

O N

O O

O

Figure 15.5

Molecular structure of 1-pyrenebutanoic acid succimidyl ester

H2N Toluene +

C2H5O Si

OC2H5

N

60°C, 24 h C2H5O Si

OC2H5

OC2H5 OC2H5

O 1-pyrenecarboxaldehyde

3-aminopropyltriethoxysailne

N-(3-triethoxysilylpropane)-1-pyrenemethaneimine

Figure 15.6 Preparation and molecular structure of N-(3-triethoxysilylpropane)-1-pyrenemethaneimine

Intensity, a.u.

2.74 nm

(E)

2.34 nm

(D)

1.96 nm

(C)

1.53 nm

(B) 1.18 nm

2

4

6 2θ / deg. CuKα

8

(A) 10

Figure 15.7 X-ray diffraction patterns of thin film of silylated graphite oxide containing perfluoroalkyl groups (A): before and after reacted with N-(3-triethoxysilylpropane)-1-pyrenemethaneimine of (B): 0.1, (C): 1.0, (D): 10 and (E): 100 mmol/L at 85 °C for 1h. (unpublished data)

[(C6F13C2H4Si(OH)2]2.3Si14O29 without amino groups were easily removed by washing with acetone. An example of the procedure shown below is the reaction of silylated graphite oxide containing perfluoroalkyl groups with pyrene containing triethoxysilane (N-(3-triethoxysilylpropane)-1-pyrenemethaneimine) prepared as shown in Figure 15.6 [69]. This reaction is basically the same as that of the first step of the procedure shown in Figure 15.4 and the silylation of silanol groups by triethoxysilyl groups. Though the molecular size of this dye is much larger than that of 3-aminopropyltriethoxysilane, the (001) X-ray diffraction peak further shifted to lower angles as the increase of the concentration of dye as shown in Fig. 15.7. The interlayer spacing increased from 1.18

478

Functionalized Inorganic Fluorides

to 2.74 nm and this strongly indicates that the dye molecules were attached to the layers of silyalted graphite oxide. This strongly indicated that the dye molecules were attached to the layers of silyalted graphite oxide.

15.3

15.3.1

Photophysical and Photochemical Properties of Dyes in Intercalation Compounds Containing Perfluoroalkyl Groups Microenvironment Estimated by using Probe Molecules Showing Photophysical Responses

The microenvironment of the microcavity in the intercalation compounds containing perfluoroalky groups available for photofunctional molecules can be estimated by using some probe molecules showing photophysical responses [70]. Polarity plays a major role in many fields and this term is used to express the complex interplay of all types of solutesolvent interactions. As mentioned in introduction, pyrene molecules are often used to estimate micropolarity. Figure 15.8 shows the fluorescence spectra of pyrene introduced in (C8F15H2NH2)0.38GO, together with that in (C8H17NH2)0.65GO. Five peaks are observed at 373–393 nm due to vibronic fine structure in both spectra and the I1/I3 ratios were 0.60 and 0.73, respectively. This indicates that the micropolarity of the interlayer space of intercalation compound of graphite oxide became lower when it contained perfluoroalky groups as expected from the lower micropolarity in fluorine-substituted methylcycrohexane [13]. In addition, pyrene is a good excimer-forming probe molecule and when pyrene molecules are aggregated, a broad peak is observed at longer wavelength than that of monomer. Control of the aggregation state of dye molecules is an important issue in artificial photosynthesis, nonlinear optics, solid dye laser, etc. In the case of layered materials, several techniques such as ‘spacer effect’ of co-intercalated surfactant molecules [71, 72] and ‘size-matching effect’ [73–75] have been reported for the suppression of

Fluorescence, a.u.

3 5 2 4 1

(B): I1/I3 = 0.60

(A): I1/I3 = 0.73 360

400

440

480 520 wavelength / nm

560

600

Figure 15.8 Fluorescence spectra of pyrene in (C7F15CH2NH2)0.38GO, together with that in (C8H17NH2)0.4GO. (unpublished data)

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups

479

Fluorescence, a.u.

aggregation of dye molecules. In the former case, when surfactant molecules such as alkyltrimethylammonium ions are co-intercalated together with dye molecules, it locates between the adjacent dyes and avoids the aggregation of them by acting as a ‘‘spacer’’. In the latter case, when the separation distances between the adjacent anionic sites in the host material and between the adjacent cationic charges in the porphyrin molecules were approximately equal, a surprisingly large value of 0.8 porphyrin molecules/nm2 for the loaded amount of porphyrin without aggregation is achieved. A similar ‘spacer effect’ of perfluoroalky groups is also expected, if hybrid intercalation compounds containing perfluoroalky groups are used as hosts [76–78]. As shown in Figure 15.9, when the

360

monomer

excimer

400

440

480 520 wavelength / nm

560

600

Figure 15.9 Fluorescence spectra of pyrene attached to silylated graphite oxide containing perfluoroalkyl groups reacted with N-(3-triethoxysilylpropane)-1-pyrenemethaneimine of (A): 10 and (B): 100 mmol/L at 85 °C for 1 h. (unpublished data)

concentration of N-(3-triethoxysilylpropane)- 1-pyrenemethaneimine was enough low, the fluorescence due to pyrene excimer around 490 nm was not observed, indicating the suppression of aggregation of pyrene attached to the film of silylated graphite oxide containing perfluoroalkyl groups. Rose Bengal shown in Figure 15.10 was also used to estimate the micropolarity in the ((((perfluoroalkanoyl)amino)ethyl)hexadecyl)dimethylammoniumion-intercalated saponite dispersed in benzene, based on its absorption and fluorescence maxima in the unaggregated state [79]. In this case, the absorption and fluorescence peaks of Rose Bengal are observed at lower wavelengths and the micropolariy was surprisingly higher than that in surfactant/clay hybrids without perfluoroalkyl groups. This was attributed to the larger amounts of adsorbed water molecules in the vicinity of ammonium part in hybrid with perfluoroalkyl groups. Quartz crystal microbalance measurements indicated ca. 4 water molecules per ammonium groups were reversibly adsorbed and desorbed. While the surfactant with perfluoroalkyl groups can take linear conformation of hydrocarbon and fluorocarbon chains, alkyl groups have a bent structure as shown in Figure 15.11. The linear conformation facilitates the access of water molecules, which results in the increase of the number of adsorbed water molecules and hydrophilic microenvironment. As a consequence, concentration of molecular oxygen became a high value of 4.5 104 M, when estimated from the triplet lifetime of RB in intercalation compound. Interestingly,

480

Functionalized Inorganic Fluorides I

I O

O

NaO I

O

Cl

NaO Cl Cl Cl

Figure 15.10

Molecular structure of rose Bengal

R’

R’

Water

RB

RB R

+ –

+ Rf

–

Figure 15.11 Schematic illustration of the orientatioal structure of RB and surfactants with and without perfluroalkyl groups in saponite clay

the adsorbed water molecules enhanced the adsorption of gaseous molecules such as hydrogen, deuterium, nitrogen and ethylene gases [80]. The weak interaction between CnF-S molecules was also demonstrated by the aggregation behaviour of antimony(V) porphyrins introduced in the above intercalation compound of saponite upon dispersion in benzene [81].

15.3.2

Photophysical Properties

The enhancement of fluorescence from rhodamine B molecules (hereafter abbreviated as RhB) introduced in silylated graphite oxide containing perfluoroalkyl groups has been reported [32, 82]. The n-hexadecylamine molecules were first intercalated into silylated graphite oxide in order to facilitate the exfoliation and then, RhB molecules were introduced by exfoliation/restacking method as schematically illustrated in Figure 15.12. Figure 15.13 shows the fluorescence spectrum of RhB in the resulting material, together with that in the sample without perfluoroalkyl groups. The fluorescence was divided by the absorption at the wavelength of 510 nm, which was measured separately. The fluorescence from RhB in intercalation compound containing perfluoroalkyl chains was much larger than that without them. The lower radiationless quenching by low vibrational C-F bonds [83] would be one of the reasons for the increase of fluorescence intensity. The rigid packing of

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups

NH2

NH2 NH2

Si

Si

481

NH2

Rf

Si

O

N

R

R

cast

R

chloroform

N+ O

N

N+

5.0 nm

R

O O OH

Rf

R

R

OH

R Rhodamine B

Rf

Si

hexane

R

Rf

Rf

R

Rf

Rf

Si

Si

NH2 NH2

n-hexadecylamine

Si

NH2 NH2

Si

NH2

Figure 15.12 Schematic illustration for the preparation of alkyamine-intercalated-silylated graphite oxide thin film containing rhodamine B molecules 35 000 (A)

Fluorescence/A510

30 000 25 000 20 000 15 000 10 000

(B)

5000 0 500

550

600 wavelength / nm

650

700

Figure 15.13 Fluorescence spectra of rhodamine B in alkylamine intercalated silylated graphite oxide (A): with and (B): without perfluoroalky groups. (Reprinted with permission from [32] Copyright (1999) Royal Society of Chemistry.)

alky or perfluoroalky chains around RhB molecules can also explain the fluorescence enhancement. This effect was first reported by Tohnai, et al. based on the enhanced fluorescence of the ammonium salts of anthracence-2,6-disulfonic acid in the solid state [84]. The rigidity of the packing suppressed the distortion of anthracene ring and accordingly radiationless decay process. In the present case, the larger volume of perfluoroalkyl groups surrounding RhB molecules is responsible for the increase of rigidity. The fluorescence intensity increased with the increase of perfluoroalky chain length as shown in Figure15.14 and the increase of n-hexadecylamine content. In both cases, the density of organic components increased, accordingly the increase of rigidity of packing, while the number of C-F or C-H groups to quench the emitted light increased. Therefore, the latter mechanism of the increased rigidity of packing seems dominant in this system.

482

Functionalized Inorganic Fluorides

Fluorescence / A510

35 000 30 000 25 000 20 000 15 000 10 000 5000 0

2

4 6 8 10 Number of carbon in perfluoroalkyl chain

12

Figure 15.14 Variation of fluorescence intensity of rhodamine B in silylated GO containing n-hexadecylamine as a function of perluoroalkyl chain length. The n-hexadecylamine/GO unit ratio was 3.0. (unpublished data)

15.3.3

Photochemical Properties

The photochemical dimerization behaviour of acenaphthylene included in perfluoroalkylamine-intercalated graphite oxide has been investigated in comparison with that in intercalation compounds without perfluoroalkyl groups [85, 86]. Upon UV irradiation (high pressure UV lamp, l > 340 nm), two types of dimers of acenaphthylene as shown in Figure 15.15 can be formed in the film of intercalation compound. It has been reported that more than 95 % and about 50 % of acenahthylene molecules in the singlet and triplet excited states, respectively provide syn-isomer [87, 88]. As the result, syn/anti ratios of the dimers are more than 19 and about 1 when they are exclusively formed from the singlet and triplet excited states, respectively. In octylamine-intercalated graphite oxide without perfluoroalkyl groups, the syn/anti ratio increased with the increase of the amount of included acenaphthylene and reached 9 [84]. As the increase of the alkyl chain length of alkyl amines, the syn/anti ratio decreased. At high concentrations, the distance between adjacent acenaphthylene molecules decreases, therefore, excited acenaphthylene molecules readily react with the adjacent molecule before converting to the triplet state. When the alkyl chain length becomes shorter a similar effect is expected. On the other hand, in (C7F15CH2NH2)0.38GO with short perfluorolaky chains, the syn/anti ratio was only 1.2

hν 2

acenaphthylene

Figure 15.15

+

λ > 340 nm

syn-dimer

anti-dimer

Photochemical dimerization of acenaphthylene

Hybrid Intercalation Compounds Containing Perfluoroalkyl Groups

483

10 9

Syn/anti ratio

8 7 6 5 4 3 2 1 0 0.2

0.4 0.6 Acenaphthylene/GO ratio

0.8

Figure 15.16 Syn/anti ratio of acenaphthylene dimer formed in alkylamine- or perfluoroalkylamine-intercalated graphite oxide as a function of acenaphthylene content. The broken lines are guides for eyes. (Reprinted with permission from [85] Copyright (1999) Royal Society of Chemistry.)

even at high acenaphthylene contents, which indicated that most of the dimers were mostly formed via triplet excited state as shown in Figure 15.16 [86]. This was ascribed to the weaker interaction between perfluoroalkyl groups and acenaphthylene molecules, facilitating the diffusion of excited acenaphthylene molecules. This allows the excited molecules in the singlet state to convert to the triplet state and this leads to the increase of the anti-isomer of acenaphthylene dimer. Photochemical decomposition in surfactant-intercalated saponite was investigated under visible light irradiation. The rate constant of the decomposition in C3F-saponite was the smallest. Photochemical decomposition of Rose Bengal is caused by the singlet oxygen, which is quenched by water molecules. The high water content around ammonium group shown in Section 15.2.4 was responsible for the decomposition [79]. Azobenzene derivatives are well known photochromic dyes which can undergo trans-cis isomerization under visible or ultraviolet irradiation as shown in Figure 15.17. Inoue et al. R1

N=N

N=N hν hν’ or Δ R2

R2

R1

Figure 15.17 Photoisomerization of azobenzene derivatives

484

Functionalized Inorganic Fluorides

reported the photochemical cis-trans isomerization of C3F-azoþ in the interlayer spacing of the scrolled nanotube of K4Nb6O17 [17]. Upon UV light (356 nm) irradiation, the absorption band at 357 nm due to trans-isomer decreased and that at 449 nm ascribed to n-p* transition of cis-isomer increased. Upon visible light (458 nm) irradiation, the absorption spectrum recovered to the starting one and this spectral change can be repeated many times. The other photoprocesses such as photoinduced electron/energy transfer and photocatalysis, which have been reported in layered clays [89], are also expected in layered materials containing perfluoroalkyl groups.

15.4

Conclusion and Future Perspectives

Hybrid intercalation compounds containing perfluoroalky groups can be obtained from various layered materials using conventional intercalation techniques including ion exchange, acid/base reaction, silylation and oxidative intercalation. Appropriate techniques should be selected, depending on the nature of the host layered materials such as ion exchange capacity, functional groups on their layers, etc and the size of the intercalated species. The resulting hybrid intercalation compounds are hydrophobic and, therefore, they are exfoliated in appropriate organic solvents, forming nanosheet solution. Integration of the nanosheets is expected to realize hierarchically controlled micro- to macroscopic structure by using self assembly, layer-by-layer and LB techniques. The microcavity surrounded by perfluoroalkyl groups available for organic molecules is generated in the hybrid intercalation materials. Various organic molecules can be introduced into the microcavity by ion exchange, exfoliation/restacking and silylation techniques. The perfluoroalky environment in the intercalation compounds offers unique effects on the introduced organic molecules in terms of low micro-polarity, higher diffusivity, higher rigidity, lower radiationless quenching, etc. These properties are ascribed to the nature of highly electronegative fluorine atoms and C-F bond with high polarizability and lower vibration energy than that of C-H. By utilizing these unique properties of hybrid intercalation compounds containing perfluoroalky groups and selecting appropriate host and guest species, various applications are expected in the fields of nonlinear optics, solid dye laser, artificial photosynthesis, recording, electroluminescence, photocatalysis and so on.

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16 The Fluoride Route: A Good Opportunity for the Preparation of 2D and 3D Inorganic Microporous Frameworks Jean-Louis Paillaud, Philippe Caullet, Jocelyne Brendle´, Ange´lique Simon-Masseron and Joe¨l Patarin Equipe Mate´riaux a` Porosite´ Controˆle´e (MPC), Institut de Sciences des Mate´riaux de Mulhouse, LRC-CNRS 7228, Ecole Nationale Supe´rieure de Chimie de Mulhouse, Universite´ de Haute Alsace, 3 rue Alfred Werner, 68093 Mulhouse Cedex, France

16.1

Introduction

The ancient mineralogists and chemists [1] identified the mineralizing role of fluorine in hydrothermal synthesis but it was only at the end of the 1970s that Flanigen and Patton [2] first used the fluoride ion for the synthesis of microporous materials with the formation of silicalite-1 in almost neutral media. Then, from 1986, after the pioneering work by Guth, Kessler and Wey at the University of Haute Alsace (Mulhouse, France) concerning the MFI-type zeolite [3], our group and then several other teams developed the fluoride route on a large scale, initially on silica-based zeolites and later on metallo- (alumino- and gallo-) phosphate microporous materials. As will be discussed later, this new F mineralizer leads, specifically, to materials with fewer defects. Thus, parallel to the work performed in Mulhouse, the fluoride method was extensively used, in particular, by Corma’s group (silica-based materials) in Valencia [4], Spain Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

490

Functionalized Inorganic Fluorides

and Ferey’s group (phosphate-based materials) in Le Mans then in Versailles [5, 6], France. The main goal of these researches on new microporous materials was the preparation of good catalysts or materials usable in new applications [7]. Throughout recent decades, this has led to the discovery of numerous products from this class by several groups around the world [8]. In Valencia, the fluoride route has led to new stable zeolites known as the ITQ-n series [4, 9, 10]. It is noteworthy that, during the same period, many new aluminosilicates of the SSZ-n family at Chevron Texaco were also discovered using the more classical alkaline route but, more recently, also the fluoride route [11, 12, 13]. In the latter cases, special attention was focused on the nature of the organic structure-directing agent (OSDA). Another interesting result was obtained with the fluoride method in combination with the use of two different sources of framework atoms, in particular germanium and silicon atoms, which made it possible to crystallize new zeolitic topologies, many of them possessing F-containing d4r (double four-membered ring) composite building units in their structures. While the templating role of the F ion in the formation of the d4r unit [14] in silicate was well established, germanium itself also favours the formation of such d4r units [15]. The spin-1/2 of the isotope 19 F nucleus is 100 % naturally abundant, which makes 19F solid state NMR spectroscopy a powerful tool for the local surrounding study of this atom inside the structures. In our group, besides the production of 3D networks (Mu-n and IM-n families), the use of the fluoride route was also successfully extended to the synthesis of clays such as beidellites, montmorillonites, saponites and new organic–inorganic hybrids having a 2:1 layered structure [16, 17]. Interestingly, for this class of 2D frameworks, the fluoride ions were proved to decrease the crystallization times and increase the crystallinity of the solids. Eight years ago we published an account on the synthesis of AlPO4s and other crystalline materials [18] then, more recently, an account on the work performed by our group focusing on both silica-based and metallo-phosphates materials prepared via the fluoride route and a review on clay mineral synthesis via this route [16]. In this chapter, we extend the review to the other teams’ results including the most recent ones and we add a paragraph on the application of this method to synthetic clays. The present chapter thus displays four successive paragraphs relative to silica- (essentially Si-, (Si,Al)-), germanium- (Ge- (Si,Ge)-), phosphate-based microporous materials and finally 2D materials.

16.2

Silica-based Microporous Materials

The structure of zeolites and related materials (including the phosphate-based materials) is characterized by a framework of linked tetrahedra, each one consisting of four O atoms surrounding a cation. This framework contains open channels and/or cages. In the case of interrupted framework, OH or F groups occupy tetrahedron apexes that are not shared with adjacent tetrahedra [19, 20]. As mentioned in the introduction, the synthesis of MFI-type zeolite (pure silica and variously substituted materials) was first experienced in our group from fluoride-containing reaction mixtures in 1986 [3]. In this latter study, the possibility to use the ‘fluoride synthesis route’ in order to crystallize zeolites from slightly acidic to slightly basic media (typically in the 5–9 pH range) was demonstrated. Our recent review [14] described specifically several data relative to the synthesis of MFI-type (and also MEL-type) zeolites. A remarkable and rather general characteristic of the fluoride route is

The Fluoride Route

491

the possibility to get large crystals, probably in relation to a smaller nucleation rate and a lower supersaturation in the liquid phase. Another important advantage of this new synthesis method, demonstrated in the case of the silicalite-1, is the very low concentration or even the absence of SiO3/2O- connectivity defects in the pure silica framework. The fluoride and tetrapropylammonium (OSDA) ions being present in very similar molar amounts in the structure of silicalite-1, the question of the role of the fluoride ion, beside the templating or pore-filling role of the organic species was immediately raised. In particular, the location of the fluoride ion in the structure of silicalite-1 was, actually for several decades, a matter of debate [21, 22] until the publication of the solid state NMR works of Fyfe et al. in 2001 [23] and the single crystal crystallographic study of Aubert et al. [24] in 2002. In the latter reference, the F- ion was located in the [415262] cage* also designed as the mel composite building unit (CBU) [25] of the structure at an interacting ˚ , leading thus to SiO4/2F units with pentacoordinated silicon distance Si-F ¼ 1.915(3) A atoms in a distorted trigonal bipyramid geometry [24]. These last characteristics concerning specifically silicalite-1 will be illustrated further in the text in a more general context including the large number of zeolites prepared up to now in fluoride media. A large number of other zeolites were then prepared in fluoride media, corresponding nowadays to several tens of different topologies reported in some recent reviews [4, 9, 10, 12, 14, 26, 27, 28]. They are listed in Table 16.1 together with new topologies recently reported. Most of the topologies were obtained as pure silica forms. Actually, among all these topologies or zeolites, only a small proportion corresponds to new structures never obtained before in the absence of fluoride from conventional reaction mixtures. Accordingly, the materials of topology AST, ATS, IFR, IHW, ISV, ITH, ITW, IWV, STT, -SVR were obtained for the first time. In addition to the materials listed in Table 16.1, several silica-based microporous materials were also synthesized using the fluoride route CJS-3[64] (certainly a clathrasil), ITQ-10 [81] (beta like) and ITQ-14 [81] (BEA-BEC overgrowth), PREFER (a lamellar precursor of ferrierite) [82], SSZ-61(Standard Oil Synthetic Zeolite-61) [12, 83], Nu-86 [12], ZSM-48 [68, 86], ITQ-28 [85], and SSZ-70 (a borosilicate) [86]. An important breakthrough emerged at the beginning of the 1990s from the laboratory of Professor Corma in Valencia, Spain, implying the use of very concentrated fluoridecontaining silicate synthesis mixtures, with H2O/SiO2 molar ratios close to 10 or less. This new approach resulted in obtaining several pure silica zeolites named ITQ-n (Instituto de Tecnologia Quimica Valencia-n), with already known or new topologies, characterized by a low framework density (FD) value, smaller than about 17 Si atoms per nm3. Among these less dense frameworks, one can mention the topologies CHA [36] (FD ¼ 15.4), BEA [32] (FD ¼ 15.6), ITE [49] (FD ¼ 16.3) or ISV [46] (FD ¼ 15.4). It must be noted here that previously the only low-density zeolites were the ones associated with high lattice substitution of Si atoms by aliovalent atoms (Al, B, Ga, Zn . . .). The easier formation of pure silica low-density zeolites in fluoride media can be interpreted with simple chemical considerations as follows. In hydroxide media, the positive charges of the cationic OSDAs can only be counterbalanced by negative SiO- groups. This obviously makes the formation of low-density frameworks difficult as it would imply a large concentration of such * The notation of the cages, i.e. [nmn’m’. . .] is usual, with m, m’ designating the number of windows delimited by n, n’ (Si or T atoms), but the CBU’s notation [25] is also given in favour of a better space representation for the reader.

492

Functionalized Inorganic Fluorides

Table 16.1 Silica-based ((Si) or (Si,Al)) microporous materials prepared via the fluoride route in the alphabetic order of their code assigned by the Structure Commission of the International Zeolite Association (SC-IZA) [25] Materials SSZ-24

Organic structure directing agent (OSDA)

N-methylsparteinium or N,N,N-trimethyl-1-adamantammonium Octadecasil quinuclidine and many others SSZ-55 [(1-(3-fluorophenyl)cyclopentyl)methyl] trimethylammonium Beta tetraethylammonium or many others Si-BEC 4,4-dimethyl-4-azoniatricyclo[5.2.2.02,6]undec-8-ene Betac 4,4-dimethyl-4-azoniatricyclo[5.2.2.02,6]undec-8-ene CIT-5 N-methylsparteinium or 2,2-diethyl-5,6dimethyl-2-azonia-bicyclo[3.2.2]nonane Chabazite N,N,N-trimethyladamantammonium SSZ-33 4,4-dimethyl-4-azoniatricyclo[5.3.2.02,6]dodec-11-ene d ethylenediamine AS-1 UTD-1F bis(pentamethylcyclopentadienyl)cobalt(III) EU-1 dibenzyl-dimethyl-ammonium or benzyl-(2fluoro-benzyl)-dimethyl-ammonium or benzyl(4-fluoro-benzyl)-dimethyl-ammonium Ferrierite linear mono- or diamines or pyridine and many others Zeolite Pd tetramethylammonium ITQ-4 N-benzylquinuclidinium or N-benzyl-1,4diazabicyclo[2.2.2]octane ITQ-32 N,N,N’,N’tetramethyldecahydrocyclobuta[1,2-c;3,4-c’] dipyrrolidinium or 4-cyclohexyl-1,1dimethylpiperazinium ITQ-7 1-butyl-1-cyclohexyl-pyrrolidinium or 1,3,3trimethylspiro[6-azoniabicyclo[3.2.1]]-octane6,10 -pyrrolidinium or 1,3,3-trimethylspiro[6azoniabicyclo[3.2.1]]octane-6,10 -piperidinium ITQ-3 1,3,3,6,6-pentamethyl-6-azoniabicyclo[3.2.1]octane or (2S,6R)-1,1-diethyl2,6-dimethyl-piperidinium ITQ-13, hexamethonium IM-7 ITQ-12 1,3,4-trimethyl-3H-imidazol-1-ium or 1,2,3trimethyl-3H-imidazol-1-ium ITQ-24 4,8-(2-methyl)-ethenobenzo[1,2-c:4,5c0 ]dipyrrolium-4-methyl-2,2,6,6-tetraethyl1,2,3,3a,4a,5,6,7,7a,8,8a-decahydro ITQ-27d dimethyldiphenylphosphonium quinuclidine Levyned ITQ-29 4-methyljulolidinium with tetramethylammonium

IZA Channel systems Ref. Codes (pore openings)a AFI

1D (12)

29

AST ATS

0D (6) 1D (12)

30 31

*BEAb 2D (12$ 12$ 12) 32 BEC 3D (12$ 12) 33 *BEAb 34 CFI

1D (14)

CHA CON

12, 35 3D (8) 36 3D (12$ 12$ 10) 12

DFT DON EUO

3D (8) 1D (14) 1D (10)

37 38 39

FER

2D (10$ 8)

GIS IFR

3D (8) 1D (12)

IHW

2D (8)

40, 41 42 43, 44 45

ISV

3D (12$ 12)

46– 48

ITE

2D (8$ 8)

9, 49

ITH

3D (10$ 9)

ITW

2D (8$ 8)

50, 51 52

IWR

3D (12$ 10$ 10) 53

IWV LEV LTA

2D (12) 2D (8) 3D (8)

54 55 56, 57

The Fluoride Route Table 16.1 Materials Silicalite-2

Silicalite-1 Mordenited MCM-35 ZSM-39d ZSM-23d ZSM-12 MCM-22 SSZ-37 CJS-2, Nonasil SSZ-50 Nu-1a RUB-10 SSZ-73 SSZ-44 Sigma-2 Sodalited SSZ-75 ITQ-9 Mu-26 SSZ-31 SSZ-23 SSZ-74 Theta-1

493

(continued) Organic structure directing agent (OSDA)

IZA Channel systems Ref. Codes (pore openings)a

tetrabutylammonium þ hexamethonium þ MEL diaminododecane or 1,1-diethyl-3,5-dimethylpiperidinium or 1-Butyl-1-cyclooctylpyrrolidinium tetrapropylammonium and many others MFI Naþ MOR hexamethyleneimine tetramethylammonium 1,3-diisopropyl-H-imidazol-1-ium 1-isopropyl-2,3-dimethyl-3H-imidazol-1-ium and others hexamethyleneimine or 3-ethyl-1,3,8,8tetramethyl-3-azonia-bicyclo[3.2.1]octane 4,8-ethenobenzo[1,2-c:4,5-c’]dipyrrolium2,2,6,6-tetramethyl-1,2,3,3a,4a,5,6,7,7a,8adecahydro bis(cyclopentadienyl)cobalt(III) 2,2-diethyl-5,7,7-trimethyl-2-azoniabicyclo[4.1.1]octane or 3-ethyl-1,3,8,8tetramethyl-3-azonia-bicyclo[3.2.1]octane tetramethylammonium or pyrrolidine or N,N,N,N0 ,N0 ,N0 -hexamethyl-N,N0 -(3-oxapentandiyl)-di-ammonium 3-ethyl-1,3,8,8-tetramethyl-3-azoniabicyclo[3.2.1]octane (2S,6R)-1,1-diethyl-2,6-dimethyl-piperidinium 1-azonia-tricyclo[3.3.3.01,5]undecane Naþ 1,10 -dimethyl-1,10 -butanediyl-bispyrrolidinium 1,3,3,6,6-pentamethyl-6-azoniabicyclo[3.2.1]octane or 1,1,2,2,6,6Hexamethyl-piperidinium (6R,10S)-6,10-dimethyl-5-azoniaspiro[4.5]decane N,N,N-trimethyl-1-adamantammonium

3D (10$ 10)

12, 58

3D (10$ 10) 3D (12$ 10)

3 59, 60 MTF 1D (8) 61 MTN 0D (6) 42 MTT 1D (10) 62 MTW 1D (12) 4, 12 MWW 2 2D (10$ 10) 12, 63 NES 2D (10$ 10) 12 NON

0D (6)

RTH

2D (8$ 8)

RUT

0D (6)

66– 68

SAS

1D (8)

69

SFF SGT SOD STI

1D (10) 0D (6) 0D (6) 2D (10$ 8)

STF

1D (10)

70 12 71 72, 73 74, 75

STF

1D (10)

76

*STOb 1D (12)

12, 77 78 79, 73 12, 80

N,N,N-trimethyl-1-adamantammonium STT -SVRe 1,10 -dimethyl-1,10 -hexanediyl-bispyrrolidinium linear mono- or di-amines or 1,3-dimethyl-4,5- TON dihydro-3H-imidazol-1-ium

2D (9$ 7) 3D (10$ 10) 1D (10)

42, 64 12, 65

a D for dimensionality of the channel system, in parentheses are the number of T atoms delimiting the pore apertures, the double arrows mean interconnected channels. b The * denotes polytypic materials. c Beta zeolite containing 85 % of polymorph B, d Aluminosilicates. e Interrupted framework.

494

Functionalized Inorganic Fluorides

connectivity defects. This is clearly not the case in fluoride media, where the positive charges of the organic cations can be equilibrated by the negative charges of the occluded fluoride anions. A decrease in the water content leads, of course, to an increase in the concentration of the soluble species and in particular of the oxofluorosilicate species, but also to an increase of the fluoride-to-oxygen ratio in the latter. Thus, a higher concentration of the incorporated fluoride is obtained, which involves a higher concentration of the cationic organic species and finally a lower FD able to accommodate more organics [4]. The general tendency mentioned above about the influence of the water content on the nature of the phase obtained was thoroughly discussed by several authors, for instance in references [4, 9, 12, 27] and is illustrated with three examples in Table 16.2.

Table 16.2 Phase selectivity for pure silica zeolites prepared under the fluoride route as a function of the water content with three different OSDAs Zeolites (FD)a

H2O/Si

OSDAs

H2O/Si

Zeolites (FD)a

Ref.

N+

N+

[73]

STI (16.7)

3.5–7

1,1'-dimethyl-1,1'-butanediylbis-pyrrolidinium

14

MTW (18.2)

+ N [4]

CHA (15.4)

3

N,N,Ntrimethyladamantammonium

7.5–10

STT (17.0)

N+ [4] BEA (15.6)

a

3.6

N-benzylquinuclidinium

4.5

IFR (17.0)

FD = framework density expressed in Si atoms per nm3

Camblor et al. [4] suggested that the phase selectivity could also occur through a kinetic control in accordance with the thermodynamic calculations by Henson et al. [87]. Indeed, the latter authors found that pure silica zeolitic materials with low FD are metastable with respect to denser phases. In reality, the phase selectivity may be in some systems quite

The Fluoride Route

495

different, with the more dense phase obtained when reducing the water content or with no selectivity at all. This may be due to specific OSDA/F/SiO2 interactions which can affect the relative stability of the composites [4]. Villaescusa et al. [9] discuss in detail the synthesis mechanism in fluoride media, focusing on the catalytic and structure-directing role of the fluoride anion. Contrary to the case of silica-based zeolites prepared via the hydroxide route where OH anions are practically never found in the as-made materials (except in the cases of hydroxycancrinite or hydroxysodalite), a general observation is that fluoride ions tend to be trapped in small cages of the as-made zeolites synthesized in fluoride media. As the location of fluoride in as-made zeolites is difficult, it was only achieved (by diffraction techniques or in some cases by solid state multinuclear NMR spectroscopy) for a restrained list of materials. Table 16.3 gathers the data available for pure siliceous zeolites with the type of cages occluding fluoride anions, the corresponding 19F MAS NMR chemical shift values and the Si-F bond lengths. The Si-F distances values are the shortest distances observed between the F and Si atoms. The interesting point is that, for all the topologies referred in Table 16.3 except for the FER and MTT topologies, the cavities occluding F display at least one four-membered-ring window (4MR) and are characterized by a relatively small size, the F being generally located in a close proximity of a 4MR. One other observation is that the O-Si-O angles around the Si atoms to which fluoride is coordinated (corresponding to the Si-F distances in Table 16.3) are distorted away from tetrahedral towards trigonal bipyramidal geometry (tbp), with a angle distortion increasing with the decrease of the Si-F bond length [26]. This observation can be related to the mechanism proposed by Villaescusa et al. [9] for the F-catalysed condensation of silica. This mechanism implies the existence of a tbp silicate species (either in hydroxide or fluoride media), the existence of which is supported by many studies mentioned by Wragg et al. [26]. Clearly, in pure silica zeolites prepared in fluoride media, the positive OSDAþ and the negative F charged species cannot be considered as an ion pair at all (the distances between the anion and the positively charged nitrogen are very large) but instead the locations of the OSDA and the F ion are determined by the constraints imposed by the framework. From a theoretical point of view, Sastre and Gale [88], developed a new force field that is able to model the location and behaviour of F in silica-based zeolites. Thus, they predicted for the SFF topology a pentacoordinated Si in a SiO4/2F environment with a trigonal bipyramid geometry based on full energy minimization of the system, including the OSDA. As a general trend, they also point out that the F location in cages is

Figure 16.1 A d4r composite building unit occluding a fluoride anion (white sphere)

496

Functionalized Inorganic Fluorides

Table 16.3 Locations of F and corresponding 19F MAS NMR chemical shifts (with respect to CFCl3) values for some pure silica as-made zeolites ˚) Si-F distance (A

Materials (structural code)

19

F MAS NMR chemical shift (ppm/CFCl3)

Cage types occupied by F(CBU’s)

Nonasil (NON) Octadecasil (AST) Si-BEC (BEC) ITQ-4 (IFR) Silicalite-1 (MFI) SSZ-23 (STT)

76 38

[4158] (non) [46] (d4r)

1.836(7) 2.632(29)

64 30

38.3 68 64

[46] (d4r) [4354] (bea) [415264] (mel)

n.da 1.914(11) 1.915(3)

33 43 24

68.4 and 56,9

[4354] (bea)

Betab Betab

58.6 and 70.3 38

[4354] (bea) [46] (d4r)

1.96(3), 1.95(2) and 1.937(14) n.d.a n.d.a

Ferrierite (FER)

56.4 or 59.5

[5661] (fer)

ITQ-3 (ITE) Chabazite ITQ-13 (ITH)

69 64 38

[4454] (rth) [4662] (d6r) [46] (d4r)

ITQ-13 (ITH)

66.5

[415264] (mel)

SSZ-55 (ATS) SSZ-55 (ATS)

c

1.987(19) 1.899(19)

75 78.1

[4264] (ats) pending into the channel [415264] (mel)d [4662] (d6r) [5464] (bik) [4156] (stf)

49b, 78 49b 9, 49b 93, 94 9 93 50, 95 50, 95 31 31

79,2

[4156] (stf)

n. d.a 2.04(8) 2.11(4) 1.867(8) or 1.744(7) 1.90(2)

38 69 62 96, 75 76

UTD-1 (DON) SSZ-73 (SAS) ZSM-23 (MTT) ITQ-9 (STF) Mu-26 (STF)

c

64

e

1.997(16) or 1.99(2) n.d.a 2.002(2) 2.632(8) and 2.631(8) 1.76(3)

Ref.

a

Not determined. Polytypic materials. c Overlapping between the 19F NMR OSDA’s and occluded fluoride signals. d Uncertain location, deduced from other structures [9]. e no experimental 19F solid state NMR spectroscopy study reported. b

determined by long-range electrostatic interactions between the OSDA cations and the fluoride ions [89]. A last comment must be made about the formation of several zeolites, i.e. AST, ITH, ITW, ISV, IWV, LTA (the synthesis of pure silica ITQ-29 (LTA) needs a multistep synthesis procedure with seeding [90]), ITQ-10 [81] and ITQ-14 [81] (the last two materials are closely related to the beta family) displaying in their framework the very small composite building d4r unit. Whatever the structure is, when the fluoride anion occupies the centre of a d4r unit (Figure 16.1), the chemical shift is the same, for all the structures i.e. 38 ppm. It is important to note that in fluoride-free synthesis media no pure silica containing d4r units could ever be prepared. Indeed the Si-O-Si angles in such a

The Fluoride Route

497

polyhedral unit are very small, which lead to high constraints, and it is precisely the incorporation of the fluoride anion that stabilizes small units such as the d4r unit. As a ˚ , see consequence of the high constraints in the d4r unit, no Si–F bonds (Si-F > 2.4 A Table 16.3) are formed because of the lack of flexibility of Si to form trigonal bipyramids [91]. It is interesting to note that in these pure silica materials the whole of the d4r units are occupied by a fluoride anion, corresponding to a full charge balance of the positive OSDA charge. This corresponds moreover to a more general observation for the other pure silica zeolites prepared in fluoride media and displaying no d4r units. Such an observation obviously means that the Si-O- defects (Q3-type connectivity defects) are not present or are very scarce in these products. This represents a very important difference with the pure or high silica materials prepared in hydroxide media, characterized by a large concentration of Q3 defects (Si-O- type but also Si-OH type). The absence or quasi-absence of defects in pure silica materials synthesized in fluoride media was first evidenced by Che´zeau et al. by 29Si MAS NMR spectroscopy in the case of silicalite-1 [92]. The high degree of perfection of the network was evidenced by the absence of the resonance at around 102 ppm/TMS (corresponding to Q3 defects) and a very good resolution of the resonances relative to the Q4 species (Si(OSi)4) of the different crystallographic sites. The same observation was made afterwards by a large number of researchers for various pure silica zeolites [4, 43, 49b]. Several reasons explain the absence of connectivity defects for materials prepared in presence of fluoride. Beside the possibility of balancing the positive OSDA charge by F ions described above, the lower pH value during the synthesis favours high protonation of the condensing silicate species and then complete condensation [92]. In addition, the effect of the stabilization of Si(OSi)3OH groups through strong hydrogen bonding with Si-Ospecies, which occurs in materials synthesized in hydroxide media, is completely negligible for zeolites prepared in fluoride media. A fourth reason is related to the effect of HF evolved during the calcination of the as-made product, which can heal the possible Q3 defects initially present in low concentration [43, 78, 92]. The first evidence of a Si-F bond, was revealed by Ko¨ller et al. [49b] on the basis of 19F-29Si CP MAS NMR spectroscopy, which enhances the intensity of the resonance corresponding to the pentacoordinated Si atom (SiO4/2F). Depending on the topology (MFI, BEA, ITE, MTW, STT, IFR) [49b], the corresponding signal at room temperature occurs in the 140 to 150 ppm/TMS range, corresponding to F atoms bound to a specific Si atom, or is broader and spread out between about 115 to 150 ppm. For the latter case, the fluoride ion exchanges between different Si atoms of the framework (note that this motion can be frozen at low temperature, 130140 K). Several other publications dealt with the evidence of these pentacoordinated SiO4/2F units by NMR spectroscopy [23, 75, 97]. In the case of assynthesized ITQ-9 [75] and Mu-26 [76] (STF) zeolites prepared in fluoride media, it was shown that the 19F, 29Si MAS NMR and 19F-29Si CP MAS NMR spectra are similar with, in particular, 16 resolved silicon T sites for both materials, one of them being pentacoordinated to four framework oxygen atoms and one fluoride atom. However, from the structural study of Mu-26 performed on high-resolution powder X-ray data, it was clearly shown that the as-made material Mu-26 is different from the as-made ITQ-9. In fact, while structure elucidation from a single crystal of ITQ-9 gave only eight silicon T sites, the structure of Mu-26 displays 16 different crystallographic silicon sites linking to the local and long range order for Mu-26.

498

Functionalized Inorganic Fluorides

As a consequence of the absence or quasi-absence of connectivity defects in pure silica zeolites prepared in fluoride media, these materials display a high hydrophobic character. This leads, for instance, to interesting competitive adsorption properties, for instance on pure silica BEA-type material in the presence of mixtures of water and toluene or water and methylcyclohexane[98]. One original application developed in our laboratory during the last ten years, consisted basically in transforming mechanical energy into interfacial energy and vice versa. The thermodynamic systems consist of a porous solid (pure silica zeolites) and the non-wetting liquid water, which were shown to be able to accumulate, restore or dissipate energy. When the pressure reaches a certain value (depending for instance on the characteristic of the porous system), intrusion of the water into the porous system occurs. When the pressure is then released the phenomenon is reversible (the water goes out from the pores), corresponding to the behaviour of a ‘molecular spring’, or is irreversible (the water remains occluded in the pores), corresponding in that case to the behaviour of a ‘molecular bumper’. Several systems were studied involving different pure silica topologies, i.e. MFI-, BEA- [99, 100, 101], CHA- [102, 103] or DDR- [104] topologies. The potential interest of the reversible system is the possibility to exert a strength (during extrusion) that is constant over a long displacement, in opposition to the ‘classical’ springs, with a strength decreasing progressively upon the displacement. Molecular simulation of the experimental pressure-volume curves were also performed for various topologies [101, 105]. To end this part, we would like to discuss the effect of the introduction of a T heteroelement, typically Al, in the synthesis mixture. Two types of effects can be observed, corresponding either to the synthesis of the same isomorphously substituted topology or to the preparation of a completely different topology [4, 8]. The isomorphous substitution of Si by Al in a given topology is possible for the majority of the pure silica zeolites synthesized in fluoride media, sometimes to a great extent, for instance in BEA zeolite, IFR, ITE, STT, CHA, STF, CIF, AST . . . The incorporation of Al in the framework is limited by the charge balance, the limit being given by the number of incorporated OSDA cations. Upon the progressive aluminium incorporation, the inorganic framework which is neutral for pure siliceous zeolite becomes negatively charged, for that reason the amount of F occluded in the aluminium containing zeolite materials is lower. It is interesting to note here that, compared to the case of a pure silica material, the crystallization of aluminosilicate analogues occurs more slowly and needs, moreover, a higher OSDA concentration in the reaction mixture. Although the incorporation of Al (or B or Ti) in the framework increases the hydrophilicity of the zeolites prepared via the fluoride route, the isomorphously substituted samples still show a far greater hydrophobicity than their counterparts synthesized in hydroxide media. The polarity of the material being very important towards the adsorption properties, the catalytic properties (selectivity and/or activity) of the zeolitic materials are also different, according to the synthesis route used (hydroxide or fluoride) as exemplified for instance by (Si,Al)-BEA [106], or (Si,Ti)-BEA [107]. The introduction of a T heteroelement may also play a structure-directing effect, changing the phase selectivity of the crystallization, as was previously reported in hydroxide media [108, 109]. In some cases, this selectivity change is simply related to the fact that the isomorphous substitution is very difficult, as is the case for the pure silica SSZ-24 (AFI) or SSZ-31 (*STO) zeolites. In other cases the isomorphous substitution is easier and

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the phase selectivity change is governed by the charge-balance effect. For instance, in the presence of benzylquinuclidinium cation (BQþ), the ITQ-4 (IFR) zeolite (1D pore system) or the BEA zeolite (3D pore system) crystallises from Al-poor or Al-rich synthesis mixture respectively [4]. Schematically, when the AlO4/2 content exceeds a certain value, corresponding to 2 BQþ per u.c. in ITQ-4, the formation of a less dense phase (here the BEA zeolite), able to accommodate a larger concentration of OSDA, is promoted. It appears thus that the effect of increasing the Al content is analogous to the effect of reducing the water content (see above). Furthermore, it appears also, that in these aluminosilicate systems, the increase of the water content still induces the formation of the denser phase ITQ-4, as already observed in pure silica systems. In opposition to the tendency evidenced in hydroxide media, where isomorphous substitution of Si by Al (or B) leads, in addition to the formation of zeolites with more multidimensional porous systems, to structures with a higher population of 4MR [109], no clear tendency is visible in this respect in the case of syntheses performed via the fluoride route. It is worth mentioning the recent and interesting discovery of the aluminosilicate AS-1 of topology DFT [37]. This aluminosilicate of chemical formula |(C2H10N2)0.5|[AlSiO2] can be synthesized with ethylenediamine as the OSDA only in fluoride media. The role of the fluoride ions is here rather tricky but in this case the addition of HF favours the diprotonation of the ethylenediamine molecules. AS-1 opens new prospects about the use of amines in fluoride media for the preparation of novel low-silica zeolites. The isomorphous substitution of Si by the tetravalent Ge element also exerts a specially marked structure-directing effect, in fact towards structures displaying d4r units, and allowed to discover a great number of new topologies, as discussed in detail in the next paragraph.

16.3

Germanium-based Microporous Materials

In the periodic table of the elements, germanium is just below silicon and both elements present similar properties and reactivities. Thus, in their pioneering work on synthetic zeolites, Barrer et al. synthesized gallo- and alumino-germanate zeolites of topologies FAU and LTA in alkaline media [110]. It was only 30 years later, in 1989 at Mulhouse, that a series of silica-based MFI-type zeolites containing germanium as framework T atoms were prepared hydrothermally in fluoride or fluoride-free media at pH varying from 6 to above 13 with tetrapropylammonium as the OSDA, the general chemical formula after calcination being Si96xGexO192 with x £ 34 [111, 112]. While framework silicon atoms are exclusively tetracoordinated in zeolites, germanium atoms can be 4, 5 and 6 coordinated at the centre of GeO4 tetrahedra, GeO5 trigonal bipyramids and GeO6 octahedra, respectively. Accordingly, in fluoride-free or fluorinated aqueous or quasi-nonaqueous media, simple amines or ammonium salts as OSDAs led to the synthesis of many open framework germanates [113]. Many of them are members of the families of products ICMM-n (Instituto de Ciencia de Materiales de Madrid-n) [114–120], ASU-n (Arizona State University-n) [121–127], UCSB-n (University of California Santa Barbara-n) [128], SU-n (Stockholm University-n) [129–135] and recently IM-n (Institut Franc˛ais du Pe´trole, Mulhouse-n) [136]. In all these compounds, the

500

Functionalized Inorganic Fluorides

germanium atoms are both tetra- and hexacoordinated to oxygen atoms, pentacoordinated Ge being less frequent as e.g. in JLG-5 [137], FDU-4 [138], ASU-14 [139], ASU-16 [124] or SU-12 [130]. All these materials display interrupted frameworks and proved to be unstable upon calcination. In these kinds of frameworks, well-defined cluster-building units containing a number of GeOn polyhedra are encountered. These building units are at the origin of the molecular building units notions and scale chemistry [140], one of the most frequently met being the GeO10 cluster also called secondary building unit SBU-10 (Figure 16.2a). The SBU-10 units may be connected directly between them as in e.g. IM-14 (Figure 16.2b) [135].

a b a)

b)

Figure 16.2 The GeO10 or SBU-10 unit present in many open-framework germanates, (a) The six tetrahedra (light grey) around the rigid central four octahedra (dark grey) cluster are connected only via two vertices, which allows a free rotation of the tetrahedra; (b) view down [001] of IM-14 [136] showing the connections between the SBU-10 units

Depending on their connectivity mode, these GeOn clusters form a great number of 3D structures from dense to very open ones [133]. Note that in some cases silicon has been incorporated in the structures as in e.g. IM-13 [141], SU-21 [134] and SU-61 [133], the Si/ Ge molar ratio being in these cases close to 0.1, 0.35 and 0.3, respectively. The first pure germanates reported with a true zeolite framework topology were ASU-7 (ASV) and ASU-9 (AST) discovered in 1998 [142]. Both structures are formed by an assembling of d4r composite building units. Then, many of the silica-based microporous materials prepared in presence of germanium held the d4r composite building unit in their structure; they are listed in Table 16.4. FOS-5, a germanate of chemical formula |(CH3)3N)6 (H2O)4.5|[Ge32O64] was prepared in the presence of DABCO in fluoride medium [143]. The structure code is BEC and corresponds to the hypothetical polymorph C of zeolite Beta [144]. Interestingly, many OSDAs allow the preparation of ITQ-17 of topology BEC in fluoride media with different Si/Ge molar ratios [145, 146]. While the structure directing role of the F-ions towards the d4r units was clearly established in the case of pure silica phases (see the preceding section), germanium itself promotes the formation of such units. A relevant example is given by Ge-ITQ-17 (BEC) synthesized in fluoride-free media with 1-methyl-4-aza-1-azonia-bicyclo[2.2.2]octane as OSDA, the presence of fluoride ions in the starting gel only accelerates the synthesis [146]. A similar

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Table 16.4 Microporous germanates or germanium-based microporous materials approved by the IZA and prepared in fluoride media [25] Materials

ASU-7, SU-10 ASU-9, SU-9 FOS-5 IM-10 IM-11, ITQ29 IM-12, ITQ15 IM-16 ITQ-13 ITQ-7 ITQ-21 ITQ-22 ITQ-24

ITQ-26 ITQ-33 ITQ-34

OSDA

Si/Ge molar ratio

Channel systems (pore openings)a

IZA codes

Ref.

dimethylamine or propylamine

0, 0.45

1D (12)

ASVb

1,4-diaza-bicyclo[2.2.2]octane or 2,2-dimethyl-propane-1,3-diamine 1,4-Diaza-bicyclo[2.2.2]octane hexamethonium Kryptofix K222 or 4methyljulolidinium with tetramethylammonium (6R,10S)-6,10-dimethyl-5-azoniaspiro[4.5] decane or 1,1,3- trimethyl6-azonia-tricyclo-[3.2.1.46,6]decane 3-ethyl-1-methyl-3H-imidazol-1ium hexamethonium

0, 2.70

0D (6)

ASTb

0 0, 0.32 £3

3D (12$ 12) 0D (6) 3D (8)

BECb UOZb LTAb

142, 163 142, 163 143 177 56, 164

4 and <10

2D (14$ 12)

UTLb

1.4

3D (10$ 8)

UOSb 165

6

3D (12$ 10$ 9)

ITHb

150

ISVb

166

1,3,3-trimethyl-6-azonium tricyclo[3.2.1.46,6]dodecane N-methylsparteinium 1,10 -dimethyl-1,10 -pentanediyl-bispyrrolidinium hexamethonium or 4,8-(2methyl)ethenobenzo[1,2-c:4,5c’]dipyrrolium-4-methyl-2,2,6,6tetraethyl-1,2,3,3a,4a,5,6,7,7a,8adecahydro 1,3-bis(triethylphosphoniummethyl)benzene hexamethonium trimethylen-1,3-bis(trimethylphosphonium) ethyleneglycol

(Si,Ge)SOD Si,Getetrapropylammonium or MFI, propylamine SU-11 SU-16 N1-(2-amino-ethyl)-ethane-1,2diamine SU-15 diisopropyl-amine SU-32 diisopropyl-amine SU-46d N1-(2-amino-ethyl)-ethane-1,2diamine

1 1.7 4 2

152, 153

b, c 147 3D IWWb 167 (12$ 10$ 8) 3D IWRb 53, 178, (12$ 10$ 10) 179

IWSb

£4

3D (12)

2 9

b, c 3D (18$ 10$ 10) 3D (10$ 9) ITRb

180, 181 169

3

0D (6)

SOD

170

£4

3D (10$ 10)

MFI

B/Ge ¼ 0.5

3D (12$ 8)

SOS

111, 112, 163 171

1.4 3D (12$ 9) 1.2 3D (10$ 8) Al/Ge ¼ 0.67 3D (8)

168

SOFb 172 STWb 172 SBN 173

(continued overleaf )

502

Functionalized Inorganic Fluorides

Table 16.4

(continued)

Materials

OSDA

SU-57d

ethylenediamine

UCSB-3d UCSB7d,e UCSB9d,e UCSB15d,e

ethylenediamine N1,N1-Bis-(2-amino-ethyl)-ethane1,2-diamine methylamine piperazine

Si/Ge molar ratio

Channel systems (pore openings)a

IZA codes

Ref.

1.85<(AlþSi)/ 3D (8$ 8$ 8) DFT Ge <19 Ga/Ge ¼ 1 3D (8$ 8$ 8) DFT Ga/Ge ¼ 0.2 3D (12) BSV

174

Ga/Ge ¼ 0.67 3D (8)

SBN

176

Ga/Ge ¼ 0.33 1D (10)

BOF

175

175 175

a D for dimensionality of the channel system, in parentheses are the number of T atoms delimiting the pore apertures, the double arrows mean interconnected channels. b Topologies containing d4r composite building units. c Code not yet assigned for ITQ-21, ITQ-33. d Synthesized in fluoride-free media with ethyleneglycol as a co-solvent. e Old material for which a code has been assigned only recently.

effect was observed for ITQ-21 [147–149], ITQ-13 (ITH) [150, 151] and ITQ-15 (UTL) [152]. On the contrary, IM-12 of topology UTL synthesized with a different OSDA than ITQ-15, namely (6R,10S)-6,10-dimethyl-5-azonia-spiro[4.5] decane requires a basic fluoride-free medium to be formed [153]. In this case, the addition of HF leads, in stirring mode, to IM-9 [154] (unsolved structure) or to (Si,Ge)-Mu-26 (STF) [76] for higher Si/Ge molar ratios at 170 C, or to Mu-31 [154b] (Mu standing for Mulhouse), when the hydrothermal treatment is performed at 150 C. Mu-31 is a molecular compound containing isolated d4r units the structure of which is close to that of [DECDMP,F]-D4R-GeO2 [155]. Note that the OSDA 4,9-dimethyldecahydro-1H,5H-dipyrrolo[1,2-a:10 ,20 -d]pyrazinediium also allows the synthesis of ITQ-25 [156] with the same topology UTL. The orienting role of germanium towards the d4r unit was extensively studied both from experimental and theoretical points of view. Accordingly, Gies et al. [15] synthesized solid solution series of AST type zeolites in the GeO2-SiO2/OSDA system where the OSDA was dimethyl-diethylammonium or dimethyl-diisopropylammonium or isopropyl-trimethylammonium cations. From quantitative 19F solid state NMR, they showed that germanium is first inserted in the d4r-unit following a Loewenstein [157] avoidance-rule. Zwijnenburg et al. [158] explained the orienting role of germanium towards d4r-containing frameworks from the destabilization energy of the d4rs compared to larger-faced cages. It was also concluded that the Ge/(Si þ Ge) ratios direct the symmetry lowering of the d4rs, which should direct the symmetry of the whole zeolite structure. From a modified model using DFT calculation, Corma et al. [159, 160] showed on clusters of chemical composition [Si8-yGeyO20H8F] (y ¼ 0, 1, 8) that Ge is pentacoordinated in the d4r-unit with a good correlation between experimental and calculated chemical shifts of the 19F NMR spectra. The authors suggested that the presence of GeO4F species could be at the origin of the fact that structures containing the d4r composite building units are more easily synthesized with germanium than in pure silica systems. Indeed, in pure silica materials, the F ion is always located at the centre of the small d4r cage with no Si-F bond (see preceding

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section). In the (Si,Ge)-system the experimental 19F chemical shift of the F occluded in the d4r unit varies between 5 and 20 ppm (with respect to CFCl3) depending of the Si/Ge molar ratio of the unit. From lattice energy minimization calculations using a special force field, Mellot-Draznieks et al. explored the known and hypothetical structures and the energetics of d4r-based zeolites with different Si/Ge molar ratios [161]. It was shown that, for the same hypothetical topology, the d4r-based GeO2 zeolites are comparatively more stable than the corresponding SiO2, AlPO4 and GaPO4 frameworks. The role of germanium in the formation of structures containing d4r-units was also investigated by Kamakoti and Barckholtz [162]. For this purpose they used a combination of DFT calculations on the periodic structure and small clusters. Their results on small clusters imply that, Ge-O-Si and Ge-O-Ge connections stabilize the d4r-units at more acute angles in contrast with the Si-O-Si linkages. Concerning the role of the OSDA in the Si,Ge-system, we shall mention the specific role of hexamethonium. Depending on the chemical composition of the starting gel, this diquat leads to four different topologies, namely the materials IM-10 (UOZ) [177], ITQ-13 (ITH) [150], ITQ-24 (IWR) [53, 178, 179], and ITQ-33 [180, 181]. Except for IM-10, these products are prepared both in fluoride or fluoride-free media. Using the high-throughput (HT) synthesis technique, Corma et al determined very complex phase diagrams and found the conditions to prepare ITQ-33, a material exhibiting straight large pore channels with circular apertures of 18-rings along the c-axis (Figure 16.3), interconnected by a bidirectional system of 10-ring channels. The resulting large micropore volume is close to 0.30 cm3g1. The HT technique was also used to synthesize new ITQ-24 (IWR) zeolite polymorphs (the first polymorph was initially prepared in the Si,Al,Ge-system) both in fluoride and fluoride-free media [53], and ITQ-30 [182, 183]. It is notable that the simple molecule diisopropyl-amine allowed the cocrystallization of the silicogermanate SU-15 (SOF) and SU-32 (STW) [172] in fluoride media. SU-32 is the fourth tetrahedral microporous material that has a chiral framework after the zincophosphate of topology CSZ NaZnPO4H2O [184], the mineral goosecreekite (GOO) and the beryllosilicate OSB-1 (OSO)[186].

b a

Figure 16.3 Projection along [001] of the ITQ-33 structure showing the 18-MRs windows

504

Functionalized Inorganic Fluorides

Finally, the ionothermal synthesis route developed since 2004 by Morris et al. [187] was not successful for the silica-based materials but the ionic liquids themselves can act as classical OSDAs in aqueous media. Thus, by using the fluoride route, Zones et al. synthesized pure silica zeolites of TON, ITW and MTT topologies using the OSDAs 1,3-dimethyl-3H-imidazol-1-ium, 1,2,3-trimethyl-3H-imidazol-1-ium and 1,3-diisopropyl-3H-imidazol-1-ium, respectively [12]. Then, in 2009, we reported the synthesis of IM-16 (UOS) a germanosilicate prepared with the ionic liquid 3-ethyl-1-methyl-3Himidazol-1-ium as OSDA in aqueous media. This microporous material possesses a new topology built from d4r and mtw composite building units [165]. The ionothermal synthesis route, up to now, was only successful in the formation of new phosphate-based microporous materials, as mentioned in the next paragraph.

16.4

Phosphate-based Microporous Materials

The fluoride route was initiated [188] by our group for the synthesis of alumino- and gallophosphates. The synthesis was similar to that described above for silica-based materials and performed by solvothermal treatment from aqueous or quasi-nonaqueous medium (ethylene glycol as main solvent for instance) of a mixture containing an aluminium (aluminophosphate) or gallium (gallophosphate) source, a phosphorous source (mainly phosphoric acid), a structure-directing agent (mainly amine derivatives) and a fluorine source (mainly HF). The first results were exciting with the synthesis of aluminophosphate AlPO4-5 (AFI-structure type) [189], the tetragonal variant of AlPO4-16 (AST structure type) [190], the triclinic form of AlPO4-34 (CHA structure type) [191] but also with the discovery of the LTA-type gallophosphate [192] and the gallophosphate cloverite (-CLO structure type) [193], which is still the inorganic molecular sieve with the largest 3-D pore channel system. Its structure displays pore openings delimited by 20 T atoms (T ¼ Ga and P) with a pore size close to 1.3 nm. Since then, several other research groups in the world developed this route for preparing new microporous phosphate-based materials. Two extensive reviews on these solids have been reported by Patarin et al. in 2002 [18] and more recently by Loiseau et al. (2007) [6]. Fluoride in these solids can compensate the charge of the organic template (ionic pair) but in most of the threedimensional solids, F is part of the framework. Three main locations can be observed: (i) F can bridge two aluminum or gallium atoms increasing their coordination from 4 to 5 or 6. More seldom, F can be tri-coordinated like in the fluorogallophosphates Mu-16 and Mu-28 where it is bound to three gallium atoms [194]; (ii) it can also be present as terminal Ga-F groups. This situation is mainly encountered for the chain- and layered solids but also for some three-dimensional solids leading to interrupted frameworks. In some cases, like for the hydroxyfluorogallophosphate Mu-5 [195], F shares this position with an hydroxyl group; (iii) In the third situation, F- is found located inside the small cubic building units, the so-called d4r units (Figure 16.4a). The presence of such a fluorine species is unambiguously evidenced by 19F MAS NMR spectroscopy since it leads to a signal with a chemical shift of about 70 ppm (with respect to CFCl3) for gallophosphates and ca. 92 ppm for aluminophosphates. Although such an environment for F was first observed

The Fluoride Route

505

for the clathrasil octadecasil [30], a large number of fluorogallophosphates, belonging mainly to the Mu-n family and showing this particularity were obtained. It is notable that in the absence of fluoride in the starting mixture, any of these materials is obtained. For instance, in the absence of F, GaPO4-a crystallizes instead of cloverite [196]. Therefore, beside its mineralizing role, F probably plays a structure-directing role stabilizing these small building units. In some rare cases, for the layered gallophosphate DMAP-GaPO [197] and for the aluminophosphate zeotype SSZ-51A of topology SFO [198], a fluorine atom is occluded in a d4r unit opened out on one edge with the GaO4 and PO4 tetrahedra (Figure 16.4b), or with the AlO4, AlO4F and PO4 polyhedra, respectively.

b a c

c

a b

Figure 16.4 Projections of two related fluorogallophosphate structures (left) Mu-5 and (right) DMAP-GaPO with the d4r-F and side opened d4r-F units, respectively (oxygen atoms of the inorganic framework have been omitted for clarity). White, light and dark grey spheres represent fluorine, phosphorus and gallium, respectively

On the other hand, F ions can be considered as catalysts. Indeed, as shown in Table 16.5 which reports some synthesis conditions of the gallophosphate Mu-32, the presence of fluorine in the synthesis mixture is necessary for its crystallization otherwise a dense phase is obtained [199]. Nevertheless no fluoride atoms were detected in this phosphate-based material. Table 16.5 Syntheses performed in the system 1 Ga2O3 : 1 P2O5 : x HF : 1 Gly : 1.5 TriPA : 160 H2O (Gly : glycine, TriPA: tripropylamine) at 170 C [199] Sample

x

A B C

1 1 0

a

Crystallization time (days) 3 7 6

XRD results Mu-32 Mu-32 þ impurity a GaPO4 quartz

Not identified, traces.

Recently fluoride ions were introduced in the starting mixtures of ionothermal syntheses in order to prepare (cobalt)aluminophosphate molecular sieves [200, 201]. In these systems, the solvent is an ionic liquid or eutectic mixture and, in many cases, can be considered as a precursor of the structure directing agent. It seems that F also plays important roles for the crystallization in these media containing few water molecules,

506

Functionalized Inorganic Fluorides

particularly on the in situ decomposition of the ionic liquid solvent [202] as it was also observed in the case of AlPO4-GIS (GIS) for which the dimethylformamide molecules are decomposed in situ to give dimethylamine as the effective OSDA [203].

16.5

Synthetic Clays

Clay minerals are defined as phyllosilicate minerals and minerals that impart plasticity to clay and that harden upon drying or heating [204]. Phyllosilicates contain two types of sheets: tetrahedral (T) and octahedral (O). The 1:1 (T:O) layer structure consists of the repetition of one tetrahedral and one octahedral sheet. The unit cell includes six octahedral and four tetrahedral sites. One surface of the layer consists entirely of basal oxygen atoms (Ob) belonging to the tetrahedral sheet (Figure 16.5a). In the 2:1 (T:O:T) layer structure, one octahedral sheet is sandwiched between two tetrahedral sheets and, in this case, six octahedral sites and eight tetrahedral sites characterize the layer unit. The tetrahedral layer sheets are inverted and two-thirds of the octahedral hydroxyl groups are replaced by tetrahedral apical oxygens (Figure 16.5b). Both surfaces consist of tetrahedral basal oxygen atoms (Ob). Structures with all the six octahedral sites occupied are called trioctahedral whereas if only two-thirds are occupied, structures are referred to as dioctahedral.

c b

a

c b (a)

a

(b)

Figure 16.5 Schematic representation of (a) a 1:1 layer and (b) a 2:1 layer present in clay minerals

The structural formula is usually reported on the basis of the half unit cell content, i.e. it is based on three octahedral sites. For example Na0.4(Al1.6Mg0.4&)Si4O10(OH)2, where & represents a vacancy in the octahedral sheet, is the chemical formula per half a unit cell of a synthetic montmorillonite, a 2:1 dioctahedral phyllosilicate [205]. These structural features, however, are limited to idealized geometric arrangements. Real structures of clay minerals contain substantial crystal strains and distortions, which lead to irregularities such as deformed octahedra and tetrahedra, ditrigonal symmetry modified from the ideal hexagonal surface symmetry and puckered surfaces instead of the flat planes made up by the basal oxygen atoms of the tetrahedral sheet. The misfit between the lateral dimension of the tetrahedral and octahedral sheets is one of the major

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causes of such distortions. The matching of the lateral dimensions can follow different mechanisms including rotation of adjacent tetrahedra, an increase in the thickness of the tetrahedral sheet or tilting of the basal oxygen plane [206]. Al3þ can replace Si4þ in the tetrahedral sheet and therefore cause a negative charge but in most phyllosilicates, this substitution is quite small. Most variation of cations is found in the octahedral sheets. The most common substitutions are Al3þ for Mg2þ and Mg2þ for Liþ. Because of the net negative charge on many phyllosilicate lattices induced by isomorphous substitutions, there are charge compensating cations between the basal oxygen planes (called the interlayer space). These cations are mainly mono and/or divalent and are surrounded by water molecules. Depending on the kind of cations and the layer surface charge, cations can easily be exchanged (sodium) or on the contrary, are tighly fixed in place (potassium). Among the 2:1 phyllosilicate group, the subgroup of smectites becomes the most interesting for modifications and applications. This is due to several specific characteristics compared to other phyllosilicates such as their high specific surface area, moderate layer charge, high degree of layer stacking disorder, large cation exchange capacity, propensity for intercalating organic compounds and ability for some of them to show extensive interlayer swelling, which can lead to a delamination of the layers. Among all the natural smectites, montmorillonite and hectorite are of particular interest because both are abundant and relatively inexpensive. The main drawbacks for their use in some applications are that natural clay minerals are mostly a mixture of several compounds, being therefore inhomogeneous both in the chemical and phase compositions. One way to circumvent this problem is to produce synthetic or semi-synthetic clays in order to control the chemical purity (and especially the absence of amorphous and gritty contaminants, as well as arsenic, iron, and other heavy metals). Beside this, their white colour assures reproducibility of brightly coloured products and the possibility of having a wide range of aspect ratio (ratio between the platelets diameters or length and thickness) opens applications in the field of polymer reinforcement [207]. Moreover, questions about clay formation and stability are best addressed by studying synthetic clay minerals that are more amenable to detailed structural characterization. Recent reviews on the preparation and properties of synthetic clays are available [16, 208–210]. In this part, we will focus on the synthesis using a source of fluoride. The route of synthesis in fluoride medium can be divided into three distinct methods: (i) semi-synthetic based on the modification of natural clay minerals, (ii) solid state synthesis starting from powders and (iii) hydrothermal synthesis.

16.5.1

Semi-Synthesis

Semi-synthetic pathways consist on the modification of natural talc minerals (having Mg3Si4O10(OH)2 as an ideal formula per unit cell) by either replacing octahedral magnesium cations by sodium or lithium [211] or by introducing metal ions in the interlayer space. For this purpose, talc is heated at high temperatures ranging from 700 to 900 C during 1 h in the presence of fluorinated salts such as Na2SiF6 or LiF (the content of fluorinated salts in the mixture can be as high as 35 wt %). It was shown that depending on the heating temperature,

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Functionalized Inorganic Fluorides

the nature of the fluoride salt and its content in the mixture, several clays can be obtained [212] like, for example, compounds having the following formula: NaMg2.5Si4O10(OH1x, Fx)2 with 0.8 < x < 1. In this case, the fluoride partially replaces the hydroxyl groups on the clay surface and thus modifies its reactivity. Expandable fluorine mica can also be prepared starting from talc and Na2SiF6 at 800 C either under air, nitrogen, argon or in vacuum. These micas seem to be formed by transformation from talc occurring without the entire disruption of the original atomic arrangement. This takes place with the loss of one magnesium cation from an octahedral site and by the intercalation of two sodium cations in the interlayer space. It was shown that a small number of hydroxyl groups in the talc structure are replaced by fluoride anions. Depending on the talc particles sizes, which are used as precursors, it is possible to tailor the particle size (from 1 to 30 mm) and therefore, the aspect ratio, which is important for applications in the fields of polymer-clay nanocomposites. Moreover, SEM micrographs reveal that beside clay aggregates, fibre-like particles are present [213]. Synthetic micas having a high charge density and containing two interlayer sodium ions, Na2Si6Al2Mg6O20F4nH2O, [214] three interlayer sodium ions, Na3Si5Al3Mg6O20F4nH2O [215] and four interlayer sodium ions, Na4Si4Al4Mg6O20F4 nH2O [216] per unit cell have been synthesized starting from talc, kaolinite or metakaolinite. Surface treatments can an also be applied in order to incorporate fluorine. Southern Clays Products Inc. commercialize TM such compounds under the trade name Laponite RD or DF containing 0.3 wt % of fluorine. These materials hydrate and swell easily to give clear and colourless colloid suspensions in water or aqueous solutions of alcohols.

16.5.2

Solid State Synthesis

This method is mainly used for the formation of synthetic sodium fluorophlogopitetype micas having either high or low charge density. Topy Industries also performed the synthesis of Li-taeniolite, Na(Mg2Li)Si4O10F2 in the same way but at a lower temperature, i.e. 700 C. Direct reaction between SiO2, MgO and Na2SiF6 in the molten state at a temperature around 1500 C leads to the formation of expandable micas having a high charge density [212a]. In this case, the aspect ratio of the materials ranges from 1000 to 5000. More recently, the preparation of low-charge mica-type clay minerals was reported [217]. Stoechiometric amounts of reactants were dry mixed in a mortar, placed in a platinum crucible and heated at temperature ranging from 700 to 950 C during 5 to 10 h in order to obtain micas of different compositions. Hydrated phases having between 0.125 and 1 interlayer sodium cations were obtained. Nevertheless, the presence of impurities such as amorphous silica and forsterite was observed in these samples. The synthesis of fluorohectorite (Na0.5Mg2.5Li0.5Si4O10F2) was also mentioned by Breu et al. [218]. They noticed that the melt without mechanical mixing yields an inhomogeneous material. Owing to the solid solution capability of the 2:1 layer silicate structure, the nonuniformity in chemical and mineral composition inevitably leads to fluorosmectites that adopt a range of charge densities. This problem was circumvented by rotating the crucible containing the mixture during synthesis which will continuously force mixing of the melt mechanically. In this case, homogeneous materials at all length scales can be synthesized.

The Fluoride Route

16.5.3

509

Hydrothermal Synthesis

The method of hydrothermal synthesis in an acidic and fluoride medium has been applied to several 2:1 phyllosilicates. The development of this route has been attempted in order to obtain dioctahedral 2:1 phyllosilicates containing Si and Al as framework forming elements under smooth synthesis conditions. Most of the clay minerals syntheses are performed in a basic medium from a gel [219–221], where the hydroxide ions are involved as mobilizing agents of the structural elements. This mineralizing role can also be played by fluoride ions, leading to the possibility to crystallize clay minerals in a neutral but also basic medium. Moreover, the substitution of F- for OH- groups in the structure introduces a local probe to gather information on the octahedral layer occupancy. Indeed, a decreasing electronegativity of the octahedral cations enhances the shielding of the fluorine nucleus (i.e causes more negative chemical shifts). It was demonstrated that there is a linear relationship between the chemical shifts and the mean electronegativity of the octahedral element [16]. The formation of 2:1 phyllosilicates is performed by using mixtures of either metal oxides or salts, water or organic solvents that provide suitable composition (mainly in stoechiometric one) for crystallization. The process is conducted at a mild temperature (a hydrothermal or solvothermal process). The hydrothermal process starts with an aqueous solution containing aluminium and/or magnesium, silicon, sodium salts that are hydrothermally treated in a stainless-steel autoclave for one to seven days to induce crystallization. Cooling and washing with distilled water and drying at low temperature complete the process. Another method is also used, which consists of preparing fine powders of suitable salts mixed in proper proportions and heating at temperatures in the range 600 to 1300 C for 2 to 10 h [222]. After cooling, the grind products are usually dispersed in water and can then be hydrothermally treated. In both procedures, a wide range of clay minerals can be obtained. The low temperature hydrothermal process is usually preferred owing to the smooth conditions. In particular, montmorillonite, with an ideal chemical formula per unit cell of Nax (Al2-xMgx)Si4O10(OH)2, where x stands for the theoretical octahedral substitution rate, can easily be obtained after 72 h of crystallization at 220 C under autogeneous pressure starting from hydrogels made of a mixture of acetates and oxides [205, 223]. It was shown that Na-montmorillonite can only be obtained in a narrow range of composition (0.2 < x < 0.4). Moreover, 19F solid-state nuclear magnetic resonance (19F MAS NMR) and extended X-ray absorption fine structure (EXAFS) reveal that there is a clustering of the octahedral element. Indeed, for a low substitution rate (i.e. x ¼ 0.2) resonances are observed at -133.3 ppm and 153.0 ppm, which correspond to Al-Al-& and Mg-Al-& fluorine environments respectively whereas, for x ¼ 0.4, only the Mg-Al-& peak is kept at 153.0 ppm, with a new one appearing at 177.0 ppm, characteristic of a Mg-Mg-Mg fluorine environment. The latter was previously observed for trioctahedral like natural hectorite and synthetic mica-montmorillonite clay. EXAFS reveals substantial increases of the bond lengths in the octahedral sheet with respect to the ideal structure. The distortions of the coordination polyhedra of divalent elements was attributed to the local trioctahedral character deduced from 19F MAS NMR spectra. In this study [205], it was shown that the substitution and the clustering of the divalent elements induce a mismatch between the

510

Functionalized Inorganic Fluorides

octahedral and the tetrahedral sheets, and finally a local undulation of the layer, which could be compensated by Al for Si substitutions in the tetrahedral sheet. This could be the reason why montmorillonite clays without tetrahedral substitutions are seldom observed in the nature. Studies carried out on the parameters governing the formation of montmorillonites in fluoride medium [205] reveal that both the crystallization time and the amount of F- in the hydrogel have significant effects on the purity of the products. In particular, 72 h of crystallization and small amounts of fluorine are necessary to obtain well crystallized montmorillonites without impurities. Commercially available fluorohectorites are also prepared starting from Na, Mg and Li silicate salts under hydrothermal conditions. In this case, products contain about 5 wt % of TM fluorine (commercially available as Laponite B ). Jacobs et al. [224] studied the influence of organic molecules on the synthesis of hectorite and stevensite from Mg(OH)2, silica sol, LiF and tetraethylammonium salt under hydrothermal conditions at 100 C for three days. They observed that both the crystallization rate and the crystallinity of these materials were increased by the addition of LiF. The addition of TEA had no effect on the rate of crystallization but improves the stacking of the layers. More recently, Carrado et al. [225] prepared a series of fluorinated hectorites starting from monovalent organic salt, LiF, Mg(OH)2, silica and magnesium chloride. After two days of crystallization, the product has a platelet size that is lower than the one usually observed for natural hectorite (19.5 nm versus 42 to 52 nm). By using organotrialcoxysilane such as phenyltriethoxysilane instead of silica, the same authors succeeded in preparing fluorohectorite in which the phenyl groups are covalently bound to the layers through Si-C bonds. This one-step functionalization avoids the ion exchange step, which is necessary to render the phyllosilicate organophilic and opens therefore new opportunities to prepare clay-polymers composites.

16.6

Conclusion

The fluoride route is at the origin of the discovery of many microporous materials. In the case of gallium phosphates, 3D networks with extra-large pore systems were found, cloverite being the main representative structure. Concerning the pure silica system, fluoride anion, in addition to its well known mineralizing agent function, seems to play a structure-orienting role towards the d4r composite building unit as for the phosphatebased materials. In this case it is located at the centre of the cubic cage. This is less true when Si is partially substituted by hetero elements. In particular, gels containing germanium as framework atom source lead to very open d4r-containing 3D frameworks, many of them also possessing extra-large pore systems, e.g. ITQ-33 for which the presence of fluoride anions accelerated only the crystallization rate and IM-12 for which fluoride is forbidden for its preparation. When the OSDAs are simple amines or ammonium salts, germanium leads, in fluoride-free or in fluorinated aqueous or quasi-nonaqueous media, to the synthesis of many open framework germanates where germanium atoms are 4-, 5- and 6- coordinated. From a theoretical point of view, it is still a challenge for the specialists to find, from computation, the accurate OSDA that chemists should use to synthesize a predicted framework or a framework existing in a different system

The Fluoride Route

511

(phosphate-based ! silica-based). Up to now, the choice of an OSDA to discover new zeolites is rather serendipitous. The fluoride route was successfully used as well for the preparation of clays free of impurities. Finally, outside of the scope of this review dealing with the preparation of 0-, 1-, 2- and 3-D crystallized microporous materials, the fluoride method was also profitably applied to produce ordered mesoporous silica materials (not discussed here) [226, 227]. This came quickly after their discovery in the early 1990s [228]. This synthetic approach was indeed advantageously used by Prouzet et al. to synthesize, in a two-step process, spherical ordered mesoporous silica with vermicular porous structure for membrane applications [229].

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17 Access to Highly Fluorinated Silica by Direct F2 Fluorination Alain Demourgues, Emilie Lataste, Etienne Durand and Alain Tressaud ICMCB-CNRS, University Bordeaux1, 87 Avenue du Dr Albert Schweitzer, 33608 Pessac cedex, France

17.1

Introduction

Silica is one of the major components of the earth and has been extensively used in both heterogeneous catalysis and chemical separations [1–5]. Numerous applications of silica are related to the hydrophobic-hydrophilic balance and surface reactivity, which both obviously depend on the surface chemistry. Because of high surface area and porous structure, mesoporous silica properties are strongly dependent on the surface reactivity of silanol species (Si-OH) and their interactions via H-bonds with other molecules such as water [6–7]. The surface of silica exhibits low Brønsted acidity without any Lewis acidity. However, surface modification reactions can be performed by using the Brønsted acidic sites of silica gels and such processes have been carried out to prepare thin coatings with a wide variety of organic groups and to adjust hydrophobic character [8]. The F-ions, generating strong ionic bonds, have been largely used to modify the surface chemistry of silica. Several studies have shown that the substitution of fluoride (F) for hydroxyl groups onto silica surface improves the acidity of the material, as illustrated by some common cracking catalysts [9, 10], as well as its hydrophobic character. As far as the synthesis routes are concerned, in nonaqueous (organic) media [11], fluoro-organosilanes have been used as precursors in sol-gel synthesis [12, 13, 14], whereas in aqueous media, alkaline fluorides such as KF, NaF or NH4F at various pH [15] are generally employed. All Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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these routes involving organic or aqueous solvents and counter ions are complex [11, 15] because the wide variety of intermediate species often leads to the decomposition of silica or SiO2 etching [16] with a subsequent formation of [SiF6]2 octahedral species [11, 15]. A recent study using aqueous medium and more especially NH4F as fluorinating agent [15] showed that very small amounts of fluorine can be chemisorbed onto silica or substituted for hydroxyls, i.e. less than 1 % in weight, because of the formation of (NH4)2SiF6. Nevertheless, most fluorinated silicas exhibit a poor thermal stability because of high reactivity with moisture [17]. Furthermore, fluorine insertion allows reduction of the electronic polarizability, the dielectric constant and the refractive index of silica [18] and allows its optical properties to be improved by increasing the transparency in visible range without parasitic optical absorption in transmission region from ultraviolet to near infrared [19, 20]. Reactive ion etching (RIE) involving CF4 or SiF4 [21, 22], as well as plasma-enhanced chemical vapour deposition (PECVD) with SF6, CF4, C2F6 , SiF4 or HF [23-27] have been developed. In the case of SiF4-PECVD, a 3.3 wt % content of fluorine could be reached. In this chapter, the reaction of elemental fluorine F2, the most reactive and the strongest oxidizing agent (Eo(F2/F) ¼ 2.87 V/ENH ) on mesoporous silica is presented [28]. Actually the half reaction ½ F2 þ e ! F is favoured by the high electron affinity of fluorine and the low binding energy (155 kJ/mol) of the F2 molecule. Such a fluorination route has been scarcely investigated because only a few laboratories possess both the equipment and expertise for research involving F2-gas [29]. SiOnFm species also provide IR and RMN reference spectra, useful for identification purposes. XPS and global titration of fluorine are generally used to determine the F/Si atomic ratios. Scanning electron microscopy (SEM) and nitrogen physisorption analysis bring information about the morphology, the nanoparticle size, the surface area and the porosity. Water, isolated silanols, as well as H-bonded hydroxyls can be identified and quantified by FTIR spectroscopy. In addition, the chemical (water uptake) and thermal stabilities of these highly fluorinated silica have been investigated for various potential applications in the field of catalysis and chemical separation. Solid state nuclear magnetic resonance (NMR) spectroscopy [30–31] has been essential to probe local environments of H, F, and Si atoms, revealing the fluorosilicate species formed.

17.2

Mesoporous Silica and Fluorination Procedures

The starting material (for instance a mesoporous silica gel n 112926-00-8, Fluka reference 60741) composed of agglomerates between 60 and 200 mm with 500 m2g1 surface area was used as-received or annealed at various temperatures under argon prior to fluorination experiments in order to remove water (T ¼ 200 C) and a part of silanol groups ( T ¼ 400 C, T ¼ 600 C). Direct F2-gas fluorination process was performed in dedicated fluorine equipment using special handling procedures previously described [29]. The sample was set in a passivated nickel boat. F2-gas was diluted in argon (Air Products). F2 volume percentages (F2 %) in the fluorinating gas were set between 10 and 80 % for reactions carried out at room temperature and between 10 and 50 % for reactions at 100 C. Results given in the following correspond to samples fluorinated during two hours at room

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temperature (RT) with F2 % in the gas mixture lower than 50 %, in order to ensure a complete reaction. At the end of the experiment, excess F2-gas was eliminated by reaction on soda lime pellets and the reaction chamber was filled with argon. It should be noted that even if the F2-gas was carefully introduced in the chamber, small amounts of SiF4 could evolve. This F2-gas route was preferred to the conventional HF liquid one because of the high reactivity of F2 (i.e. weak binding energy DHF2 ¼ 155 kJ mol1). Conversely, HF exhibits a higher binding energy (DHHF ¼ 565 kJ mol1) and, furthermore, easily decomposes silica into SiF4 gas by an etching phenomena (SiO2 þ 4HF ! SiF4 þ 2H2O). In order to limit an increase of temperature during the exothermic fluorination process, fluorine can be diluted in a neutral argon carrier.

17.3

About the Chemical Composition and Morphology of Highly Fluorinated Silica

Fluorine was quantified by the Seel method [32]. The sample was first dissolved in K2CO3/ Na2CO3 mixture by heating in a platinum boat. The molten solution was cooled to room temperature and dissolved in a small amount of distilled water. About 1 g of silica and 20 mL of 98 % H2SO4 were slowly added to the solution. This solution was then distilled under a water vapour flow at 250 C in order to favour the formation of H2SiF6 and its evaporation. H2SiF6 was condensed and the fluorine content was determined with a fluoride ion-specific electrode. XPS was used to determine the amounts of fluorine (F1s), silicon (Si2p), oxygen (O1s) and carbon (C1s) elements fixed onto the silica surface (analysis depth ¼ 5–10 nm) [28]. The analysis of the binding energies coupled to a fitting of the peaks could not allow identifying the signal of the Si-F chemical bonds [12, 33]. Infrared spectra were acquired in aim (i) to identify the different bonds present in fluorinated silica, (ii) to quantify the water and hydroxyl groups contents, (iii) to characterize the accessibility of Si-OH groups and (iv) to estimate the thermal stability of hydroxyl groups and water trapped into the network [35]. Surface-grafted fluorine content has been evaluated by XPS and for the global rate by elemental analysis (the latter being called F-titration throughout this section). The F2 percentage in the fluorinating mixture induces significant changes in the grafted fluorine content. The results of fluorinations conducted at 25 and 100 C are given in Table 17.1 and illustrated in Figure 17.1 as F/Si atomic ratios calculated from the titration results. In Table 17.1 and following sections, all samples have been noted with code number already proposed in ref. 28. As observed in most gas-solid reactions, the grafted fluorine content is obviously higher in the superficial layer of primary nanoparticles than in the bulk. The systematic difference between XPS and F-titration values observed for low F2 contents in the fluorination gas (F2 % < 50 %) corresponds to a ratio (F/Si)bulk/(F/Si)XPS close to 0.7. However, considering the small particle size (10–20 nm diameter) and the analysis depth of the XPS technique (5–10 nm), one should have to notice that the XPS analysis accounts for averaged values intermediate between the hydroxyl-rich surface and the core of a nanoparticle. As expected, the fluorine grafting increases with increasing F2 content in the fluorinating mixture. A rapid ascent below F2 % ¼ 10 % is followed by a slight increase both in XPS and total fluorine content up to F2 % ¼ 50 %. For samples fluorinated at 25 C, a gap in the F-content is noticed between 50 %

522

Functionalized Inorganic Fluorides

Table 17.1 Synthesis conditions and elemental analysis of fluorinated silica. Influence of the F2 dilution and fluorination temperature on chemical compositions of surface (XPS) and bulk (F titration by specific electrode). The fluorination duration is two hours for all samples Sample notation

Experimental conditions Fluorination temperature (C)

SiO2 10F2-25 20F2-25 30F2-25 40F2-25 50F2-25 60F2-25 70F2-25 80F2-25 10F2-100 20F2-100 30F2-100 40F2-100 50F2-100

Chemical composition measured by

F2 (vol.%)

XPS

Bulk F titration

F (at.%) F/Si F/Si F (wt.%) F/Si (–0.5) (–0.04) (–0.2) (–0.02) (–0.01)

(F/Si)bulk/ (F/Si)XPS ratio (–0.05)

25

0 10 20 30 40 50 60 70 80

0.0 8.9 9.8 11.5 13.3 12.1 13.7 14.1 14.3

0.00 0.30 0.33 0.37 0.41 0.41 0.48 0.48 0.48

2.1 2.0 1.9 1.8 1.6 1.9 1.9 1.8 1.8

0.00 3.64 6.81 7.67 8.46 7.99 12.69 11.71 12.95

0.00 0.13 0.24 0.26 0.29 0.27 0.44 0.40 0.45

– 0.43 0.73 0.70 0.71 0.66 0.92 0.83 0.94

100

10 20 30 40 50

10.1 12.5 13.8 14.3 15.2

0.34 0.42 0.48 0.48 0.51

2.0 1.9 1.9 1.8 1.8

7.54 10.99 9.67 9.70 10.51

0.25 0.38 0.33 0.33 0.36

0.74 0.90 0.69 0.69 0.71

0.6

F / Si atomic ratio

0.5 0.4 0.3 0.2 0.1 0 0

20

40

60

80

100

F2 %

&

Figure 17.1 Comparison of surface (~F2-100C XPS, ^F2-25C XPS) and bulk (nF2-100C global F-titration, F2-25C global F-titration) F/Si atomic ratios in fluorinated silica versus F2 dilution. The fluorination has been conducted during two hours at room temperature for all samples. Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

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and 60 % more clearly in the case of F-titration than for XPS analysis. The two curves tend to merge to a saturation F-content (F/Si ¼ about 0.45) for highest F2 contents in the reaction gas, which corresponds to a deeper fluorination of the agglomerates. Because of the fair reproducibility of the results, the fluorine content can be controlled with F at a percentage varying between 0 and 15 %, i.e. for F/Si values between 0 and 0.5. The rather close values obtained for higher fluorination levels on both surface and bulk, bring to the fore that the fluorination proceeds thoroughly within the agglomerates. Finally the fair agreement between the variation of F/Si atomic ratios determined by XPS (surface) and F-titration (bulk) indicates that F2 direct fluorination is homogeneous. The reaction takes place rapidly and no kinetic limitation has to be considered in this duration range. The key role of hydroxyl groups has been established by modulating their number on the surface by annealing the silica at various temperatures: 200 C, 400 C and 600 C under argon atmosphere prior to the fluorination treatment. A significant reduction of the F/Si ratio when an annealing has been performed above T ¼ 200 C, which corresponds to the beginning of dehydroxylation. For annealing at T > 200 C, a significant elimination of hydroxyl groups occurs in starting silica without modification of the texture [39] and a decrease in fluorination level takes place with a drastic diminution of both surface area and pore volume. These observations confirm that the fluorination reaction can be directly connected to the physisorbed water content as well as to the different types of Si-OH chemical bonds present in the starting material.

17.4 17.4.1

FTIR Analysis About the Content and Nature of OH/Water Groups in Highly Fluorinated Silica

Infrared spectra of starting silica before and after evacuation under vacuum at 150 C are presented in Figure 17.2. Such a (OH) pattern is characteristic of silica having a heavily hydroxylated surface [35, 36]. The sharp (OH) band at 3737 cm1 corresponds to isolated Si-OH whereas the broad band centred at about 3550 cm1 reveals the silanol H-bounded in a chain [37]. Two others minor features evidence (OH) bands of terminal (3715 cm1) and internal (3660 cm1) Si-OH groups. The latter species is inaccessible to D2O because it is located in small cavities or in narrow interspaces between grains. A schematic description of species is drawn in Figure 17.2. Infrared spectroscopy cannot differentiate isolated and geminated hydroxyl groups. Valuable information regarding these latter species will be described in Section 17.6 using the MAS-NMR technique. FTIR spectra of the pristine and silicas fluorinated at RT for two hours using F2 percentage up to 80 % have been recorded at room temperature in the 600–5500 cm1 range. In the IR spectrum of the starting silica (Figure 17.2), two characteristic bands of water at about 1630 and 5270 cm1 are assigned to d(H2O) and þd(H2O) respectively. Knowledge of the integrated molar absorption coefficients of the þd (H2O) band (1.53 cm mmol1) [38] allowed the water concentration to be determined for all samples: values are given in Table 17.2. Whereas starting silica contains about 8 wt % water, this amount rapidly decreases upon fluorination. Taking into account the specific surface

524

Functionalized Inorganic Fluorides isolated

terminal

H-bounded

internal

3737 cm–1

3715 cm–1

3550 cm–1

3660 cm–1

H

H -O

O

O-H….O-H...O-H

O-H….O-H...O-H

Si

Si

Si

Si Si - O H

Si

Si

Si

νOH

(ν + δ) H2O 5270

Si

3715 3737 0.1 (ν + δ)OH 4520

5400 5000 4600 4200

3660 3550 0.5

3500

(cm–1)

3000

δ H2O 1630

2500

2000

1500

(cm–1)

Figure 17.2 (a) Infrared spectra (transmission mode) of pristine silica in the 1400-6000 cm1 range; dotted line, spectrum recorded at room atmosphere; full line, spectrum recorded after outgassing the sample under vacuum during 3 hours at 100 C. Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

area, the concentration of molecules at the surface (nH2O per nm2) can be calculated (Table 17.3). The decrease is less abrupt and leads to conclude that the fluorination improves the hydrophobic character of silica. Table 17.2 Determination by infrared spectroscopy of adsorbed water in pristine silica and in samples fluorinated at room temperature with various F2/Ar molar ratios. Fluorination conditions as in Table 17.1 Sample

Bulk H2O (wt.%)

n H2O (mmol g1)

n H2O per nm2

SiO2 10F2-25 20F2-25 50F2-25 60F2-25 70F2-25 80F2-25

8.0 2.9 1.0 1.2 0.8 1.2 1.7

4800 1650 550 650 450 650 950

5.7 2.3 0.9 1.3 0.8 1.4 2.1

Figure 17.3 shows IR spectra for dried silica after outgassing under vacuum of 104 Pa at 150 C for 3 hours. The intensity of the (OH) band of the isolated Si-OH group at 3737 cm1 becomes weaker upon fluorination, showing the lower concentration of isolated hydroxyl groups. For the highest F contents the (OH) band of isolated silanol group finally vanishes whereas a small part of the one of H-bounded species (3660 cm1) persists. Quantification of the dependence of the residual content of surface hydroxyl groups upon the F2 percentage was obtained from the area of the characteristic þd (OH) band of silanol

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Table 17.3 Determination of silanol content from infrared data; C and n are considered as concentration and sites number respectively. Calculations are based on the following chemical formula: SiO2-x-y(OH)2xF2y. Fluorination conditions as in Table 17.1 Sample

SiO2 10F2-25 20F2-25 50F2-25 60F2-25 70F2-25 80F2-25

OHtotal COH total mmol per Si per g of dry silica 6550 4150 2800 2650 1400 2050 2000

0.42 0.27 0.18 0.17 0.09 0.13 0.13

OHacc nOH acc. nOH total COH per nm2 acc.D2O per Si D2O per nm2 of of silica mmol per g of dry silica silica 7.8 5.7 4.6 5.3 2.5 4.3 4.4

5150 2750 1500 1350 400 950 850

0.33 0.18 0.10 0.09 0.03 0.06 0.06

6.1 3.8 2.4 2.7 0.7 2.0 1.9

Fraction OHinternal of SiOH per Si acc. D2O 0.78 0.67 0.53 0.52 0.27 0.47 0.42

0.09 0.09 0.09 0.08 0.07 0.07 0.08

situated in the 4520–4560 cm1 range (Figures 17.2 and 17.3) because its integrated intensity is less sensitive to the type of silanol than the one corresponding to the (OH) band [40]. An integrated molar absorption coefficient "þd(OH) equal to 0.16 cm mmol1 is taken from [38]. For such a band located at high wave numbers, deviation from the Beer-Lambert law was taken into account through evaluation of an enhancement factor due to high diffusion of these samples. [40–42]. The silanol contents are reported in Table 17.3 together with the concentration (C in mmol per gram of dry silica) and the number (n per nm2 of dry surface) of OH groups. The concentration of OH groups varies in the following order: SiO2 > 10F2-25 > 20F2-25 > 50F2-25 > 60F2-25. We can conclude that the higher the F content, the lower the OH content. However, for F2 percentages in the treating gas higher than 20 %, the influence of the concentration of fluorine in the gas phase becomes limited. Then the formation of fluorinated surface, which limits the diffusion of reactive gas phase through the outer layer as well as the occurrence of residual Si-OH unaffected by fluorination, can be considered. Differentiation between accessible and inaccessible OH can be performed by hydrogen-deuterium (H/D) exchange with D2O. This process is followed by IR spectroscopy thanks to the downward shift of the stretching mode from 3750–3500 cm1 (OH) to 2760–2600 cm1 (OD). At first, all accessible silanols are exchanged with heavy water vapour several times (contacting with reagent excess, followed by evacuation in vacuum). After each exchange, an IR spectrum is recorded and the process is considered to be completed when there is no difference between two consecutive spectra. Modification of IR spectra for H/D exchanged samples by D2O treatment is illustrated in Figure 17.3 and the quantification of OH species is given in Table 17.3. Substituted pristine silica shows a strong decrease of intensity in the OH range and a subsequent formation of a OD band in the 2750–2600 cm1 range. In addition, the similar profiles of the (OH) and (OD) bands before and after H/D exchange respectively indicate that no supplementary hydroxyl groups are formed by the heavy water treatment. After D2O treatment, a broad maximum centred at 3660 cm1 remains. Such a band reveals unaffected internal Si-OH groups located in the interparticle space or at inner surfaces.

526

Functionalized Inorganic Fluorides SiO2

10F2-25 20F2-25 50F2-25

0.01

60F2-25

4700

4600

3737

4500 4400 (cm–1)

4300

3665 3550 2756

SiO2

0.4

2621

3660

10F2-25

20F2-25

50F2-25 60F2-25

3500

3000

2500

(cm–1)

Figure 17.3 IR spectra of pristine silica and silicas fluorinated at room temperature for different F2/Ar gas ratios; full line: spectra recorded after outgassing the sample under vacuum during 3 hours at 150C; dotted line: spectra recorded after H/D exchange by D2O treatment. Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

Furthermore, the intensity of the 3660 cm1 band of the internal Si-OH is not affected by F2 treatment. The content of internal Si-OH remains constant at about 1600 mmol g1 and the OH/Si ratio remains equal to 0.09, whatever the fluorination treatment at room temperature. 17.4.2

FTIR Bands Related to Si-F Bonds

Several significant bands of low intensity appearing in the 600–1400 cm1 range are illustrated in Figure 17.4. On the pristine spectra, the band at 975 cm1 corresponds to a Si-OH vibration mode and the bands around 1100 and 800 cm1 are related to two Si-O-Si stretching modes of

Access to Highly Fluorinated Silica

527

silica network. On the fluorinated materials, the pristine Si-OH band diminishes and two absorption bands already reported in the literature [11] grow at its expanse. The 980 cm1 band has been tentatively attributed to O3/2SiOH, O2/2SiF2 and O1/2SiF3 tetrahedral species. The band at 935 cm1 indicates the replacement of O3/2SiOH species by O3/2SiF ones. Moreover a band at 746 cm1 appears immediately at the lowest F content. Controversial interpretations have been proposed [11]. Most authors mention that this band is due to octahedral species with general formula [F6-nSi(OH)n]2- but the (29Si,19F) MAS-NMR investigations do not reveal this kind of species, which are indeed unstable on mesoporous silica treated under F2-gas [71]. Then taking into account the nearest Si-O-Si stretching mode band, the new band at 746 cm1 could be due to O4/2SiF species: i.e. intermediate between tetrahedral and octahedral entities or could correspond to another Si-O-Si stretching mode for which the bond is relaxed because the occurrence of strongly polarized Si-F bonds in its vicinity induces a decrease of the wave number. 1100

935

0.1 980

746 SiO22 10F2 25 10F2-25 20F2 20F2-25 50F2 50F2-25 60F2 60F2-25 70F2 70F2-25 80F2 80F2-25

1400

1200

1000

800

600

Figure 17.4 Range of the Si-F bonding in the infrared spectra [1 mg of silica sample diluted into 100 mg of KBr]. Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

17.4.3

Correlation between Silanol Groups on Mesoporous Silica and Grafted Fluorine on Highly Fluorinated Silica

Figure 17.5 clearly illustrates the relationship between the decrease of silanol content and the increase of the grafted fluorine content obtained from bulk F-titration [28]. At low

528

Functionalized Inorganic Fluorides

SiOH content / (mmol g–1)

fluorine contents a linear relationship between silanol consumption and the amount of grafted fluorine is observed. Therefore, direct replacement of hydroxyl groups by fluoride ones takes place at the beginning of the fluorination and the key role of hydroxyl groups in the fluorination mechanism is then evidenced. At high fluorine contents, the amount of grafted fluorine reaches an upper value about 6500 mmol g1 while, as noted previously, internal silanols remain unreacted. This suggests that for high F2-gas concentration a second reaction path involving the opening of Si-O-Si siloxane bridges may occur.

6

4

2

0 0

2 4 6 Bulk fluorine content (mmol g–1)

&

Figure 17.5 Variation of the total (^) and accessible ( ) silanol contents (deduced from IR experiments) vs. the bulk grafted fluorine content (quantified by global titration) for various F2 dilution in the fluorination gas. Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

Estimated chemical compositions SiO2xy(OH)2xF2y of fluorinated silica are given in Table 17.4 from F/Si and OH/Si atomic ratios deduced for each sample from F-titration and FTIR data respectively. Such formulations clearly account for the occurrence of higher rates of grafted fluorine for lower hydroxyl group contents. It should be pointed out that the amount of grafted fluorine is very high, when compared with conventional methods, ranging from 3.6 to 12.7 wt. % for F2/Ar ratios varying from 10/90 to 60/40. The high Table 17.4 Chemical compositions of bulk pristine and some fluorinated silicas as deduced from FTIR and F-titration data. Fluorination conditions as in Table 17.1 Sample

Chemical compositions SiO2xy(OH)2xF2y

SiO2 10F2-25 20F2-25 50F2-25 60F2-25 70F2-25 80F2-25

SiO1.79(OH)0.42F0.00 SiO1.80(OH)0.27F0.12 SiO1.79(OH)0.18F0.23 SiO1.78(OH)0.17F0.27 SiO1.74(OH)0.09F0.43 SiO1.73(OH)0.13F0.40 SiO1.71(OH)0.13F0.44

Bulk fluorine F- titration (wt.%) 0.00 3.64 6.81 7.99 12.69 11.71 12.95

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529

0.5

10

0.4

8

0.3

6

0.2

4

0.1

2

0

area of the 935 cm–1 band (a;u.)

F/Si ratio from titration data

fluorination level can be obviously correlated with the initial high content of hydroxyl due to the nanoparticles morphology of the used silica. The area of the Si-F band at 935 cm1 allows the increasing number of Si-F groups in the fluorinated silica to be followed. In Figure 17.6, this area and the F/Si obtained from F-titration are drawn versus the F2 percentage in the fluorination gas. Both follow a similar trend, exhibiting a knee at nearby F2 % ¼ 50.

0 0

20

40

60

80

F2 %

Figure 17.6 Variation of the F/Si ratio obtained from the global F-titration (^) and of the area of the IR 935 cm1 band taken as an indicator of the Si-F bond content (n) versus the F2 concentration in the fluorination gas. The drawn dotted line is a naive picture of the occurrence of the reaction path associated to the redox process involving F2/F and O2/OH couples. Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

Considering the F/Si atomic ratio determined by F-titration, one should point out that for F2/Ar ratios lower than 50/50, the F þ OH sum always lies around 0.4. For instance in 10F2-25 conditions, the sum is equal to 0.41, thus confirming the partial substitution of OH groups by fluoride. For higher F2 percentages, the fact that this sum is superior to 0.5 also confirms that siloxane bridges have been partially opened and that one Si-O-Si bridge has been replaced by two terminal Si-F bonds. The final applications of silica depend on the morphology, structure and chemical bonding of the pristine material. High resolution SEM analysis of the pristine silica reveals the occurrence of nanoparticles around 10–20 nm in diameter, which are packed in uncircular agglomerates of about 100–200 mm. It should be noted that the 500 m2g1 surface area specified by the supplier is in agreement with the measured SBET. Scanning electron microscopy (SEM) images are given in Figure 17.7 for various F2 %. Within the SEM resolution range the morphology of fluorinated silica appears identical to the one of pristine silica. The 10–20 nm nanoparticles constituting the agglomerates are still present after fluorination. Nitrogen physisorption can give information about the surface area and pore-size distribution of samples fluorinated at 25 C for two hours. The pore size distribution is obtained through the Barrett-Joyner-Halenda (BJH) method using the adsorption branch [34]. No change in the isotherm type is observed after fluorination, proving that the mesopore system is conserved. Fluorination of a pristine silica with SBET ¼ 508 m2g1

530

Functionalized Inorganic Fluorides

a)

b)

c)

d)

Figure 17.7 SEM photographs of pristine silica (a) and silica fluorinated at room temperature, during two hours and with various F2 dilution: 10% (b), 20% (c) and 50% (d) Reproduced by permission J. Phys. Chem. C 2008, 112, pp. 1094310951, American Chemical Society

leads to a pronounced decrease in the specific surface area while the fluorine content increases, i.e. down to 273 m2g1 for the 80F2-25 sample. The prominent feature is the decrease of the volume of the smallest pores with intensive fluorination. We can assume that the pore size measured by nitrogen physisorption corresponds to the intergrain distances between the nanoparticles forming agglomerates imaged by SEM. Weight loss corresponding to water departure during the degassing at 150 C prior to the isotherm acquisition is drastically lowered as the F2 content of the gas mixture increases. Indeed the fluorination yields a decrease in the physisorbed water amount. Whereas starting silica contains 8 wt. % water, for fluorination carried out with an F2 percentage superior to 20 %, this amount rapidly decreases down to around 1 wt. %.

17.5

Thermal Stability and Water Affinity of Highly Fluorinated Silica

The thermal stability of the fluorinated silica samples can be determined using a thermogravimetric analysis (TGA) coupled with a mass spectrometry. These techniques allow the determination of the various species released at critical temperatures and the identification

Access to Highly Fluorinated Silica

531

of the hydroxyl sites concerned by fluorination. Figure 17.8 illustrates the weight loss of fluorinated compounds measured by TGA during a heating cycle from 25 C to 900 C under argon at 5 C/min. Figures 17.9a and 17.9b illustrate the ion current corresponding to masses m ¼ 18, m ¼ 19, m ¼ 85, m ¼ 86 and m ¼ 104 correlated to the departure of H2O, F and SiF4 (three fragments) respectively [44]. The comparison of these TGA curves with that of initial silica shows that fluorinated silica contains less water than pristine silica (weak loss between 25 and 100 C), proving that the quantity of water physically absorbed is substantially lowered by fluorine introduction. The hydroxyl loss between 100 C and 250 C is almost the same for starting silica and fluorinated silicas. The first effect of fluorination appears at 250 C when the curve of silica fluorinated with an F2/Ar > 20 % gas mixture, exhibits a strong weight loss, correlated to the departure of fluorinated fragments in the mass spectrometry analysis. At 450 C this phenomenon is slightly increased and results in a higher weight loss in TGA curve and a higher ion current on mass spectrometry curve. Temperature (°C) 0

100 200 300 400 500 600 700 800 900

0

Weight loss (%)

–2

–4 10F2 SiO2

–6

20F2 –8

–10

50F2

30F2 40F2

Figure 17.8 Thermal stability of pristine and various fluorinated silica measured by TGA up to 900 C. Reproduced by permission J. Phys. Chem. C 2009, 113, pp. 1865218660, American Chemical Society

F2 treatment leads to a significant fluorination, which makes it possible to clean the surface of Si-OH groups that are usually stable, at least for low temperatures. All types of hydroxyl seem to be involved in the reaction of fluorination. Furthermore for the 10F2 sample, between 100 C and 550 C, the TGA curve exhibits exactly the same weight loss than the unfluorinated silica accounting that fluorinated group are stable up to fairly high temperature with less than 3 % weight loss at 550 C. From 550 C an important weight loss occurs followed by a departure of fluorinated fragments mostly SiF4 [45], the same as observed at 250 C and 450 C for highly fluorinated samples. Previous studies conducted on fluorinated silica obtained via various synthesis routes pointed out that these materials exhibit some drawbacks like strong sensitivity to moisture

532

Functionalized Inorganic Fluorides (a)

0

1,E-07 TG

–1 m18 'H2O'

–2

1,E-08

–5

1,E-09

m85 'SiF3'

–6

1,E-10 m86 'HSiF3'

–7 –8

1,E-11 m104 'SiF4'

–9 –10

0

100

200

(b)

300 400 500 600 Temperature (°C)

700

800

1,E-12 900

1,E-08

0 TG

–1 –2

Weight loss (%)

Ion current (A)

m19 'F'

–4

m18 'H2O'

m19 'F'

1,E-09

–3 –4 –5

m85 'SiF3'

1,E-10

m86 'HSiF3'

1,E-11

–6

Ion current (A)

Weight loss (%)

–3

–7 –8

m104 'SiF4'

–9 –10

0

100

200

300 400 500 600 Temperature (°C)

700

800

1,E-12 900

Figure 17.9 Thermal stability of fluorinated silicas followed by mass spectrometry for (a) 10F2 and (b) 50F2 samples. Ion current curves for m/z 18, 19, 85, 86, 104 correspond to the departure of water, fluorine and various SiF4-based species. Reproduced by permission J. Phys. Chem. C 2009, 113, pp. 1865218660, American Chemical Society

and weak durability and thermal stability. This reactivity was accompanied by a departure of fluorine due to the formation of fluorhydric acid. The determination of the fluorine content in fluorinated silica stored up to 50 days in an atmosphere containing 75 % of humidity (saturated NaCl) has shown that surprisingly, the resulting fluorine contents are identical to the initial ones. Moreover, the 10F2 sample exhibits the same F/Si ratio after being stored in laboratory conditions up to six months after synthesis or after five days

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533

under ultra high vacuum in the XPS chamber. These very attractive features illustrate the high fluorine stability in silica fluorinated by direct F2-gas, even in extreme conditions (moisture, ultra high vacuum). The hydrophobic character of fluorinated silica was estimated by water adsorption a 75 % humidity atmosphere [43]. After 50 days in such atmosphere all fluorinated silicas have a drastically lower water adsorption (between 2 and 13 %) than unfluorinated silica (around 30 %) [28]. This result can easily be correlated with the lower hydroxyl content in fluorinated silica, these groups being at the origin of the surface water adsorption. However the phenomenon appears more complex when the content of grafted fluorine is considered. Indeed the higher the fluorine content, the higher the water adsorption. For example 20F2 absorbs around 8 % of water, whereas 60F2 holds around 13 % of water. Moreover concerning the 10F2 sample, after a weight increase of 5 % up to the seventh day (170 h), a weight loss starts until the end of the measurement (1200 h). This unexpected phenomenon could be explained by the coexistence of two mechanisms involving the water adsorption due to the relative humidity on one hand and the weight loss due to the departure of HF on the other hand. This removal could be compared with the mechanism of HF/H2O chemical etching of SiO2 [46]. Actually, as far as the F2 treatments with F2 contents higher than 10 % are concerned, even if these samples contain a higher F amount and a lower OH rate than the 10F2 sample, the uptake due to water is always larger than the HF departure, leading to a monotonous variation of global uptake with time. In order to understand the effect of the surface groups (OH, F, water) on the water affinity, the weight uptake of F-silica pretreated at various temperatures: 200 C, 400 C and 600 C prior to fluorination treatment has been considered. Concerning the dependence of the water uptake of fluorinated samples with these annealing temperatures, a decrease is observed for increasing temperatures without changing the shape of the curve; after 50 days exposure to humidity, the weight uptake follows the sequence: 600 C<400 C<200 C, which is in agreement with the phenomena already described above, that is the reduction of surface area as a function of annealing temperature. Up to 200 C the removal of the water from the silica surface is the main phenomenon; after this dehydration step, the dehydroxylation becomes the most important process of the sample annealing, reducing the amount of hydroxyl groups content at the silica surface and therefore limiting the fluorine grafting.

17.6

Nuclear Magnetic Resonance (NMR) Investigations

Solid state nuclear magnetic resonance (NMR) spectroscopy is essential to investigate pristine silica and the changes that occur upon fluorination. For nuclear spins ½ like 1H, 19F and 29Si the species are identified through the position of their NMR signals on the spectrum. Nevertheless, crosspolarization (CP) [30–31] is routinely used in NMR of spin I ¼ ½ as a mean of enhancing sensitivity by magnetization transfer between dipolar coupled abundant (1H and 19F) and diluted nuclei (29Si). The CP technique is furthermore a valuable spectral editing tool that can be used to determine heteronuclear spatial proximities. Therefore 1H, 19F and 29Si magic angle spinning (MAS) and 1H-29Si and 19F-29Si CP-MAS NMR experiments can be performed to probe the local environments of H, F and Si atoms.

534

17.6.1

Functionalized Inorganic Fluorides

Local Environments in Highly Fluorinated Silica through NMR Experiments

Surface hydroxyl groups can be classified according to their coordination (single O3/2SiOH and geminal O2/2Si(OH)2 silanols) and their hydrogen bonding (isolated and bound silanols, the latter divided between silanols bridged to other silanols and silanols hydrogen bonded to physisorbed water). Despite severe line broadening due to a combination of strong homonuclear dipolar coupling and small chemical shift dispersion, previous 1H NMR studies of silicas allowed identification of the resonances of water, bound silanols, and isolated silanols. In the literature, the 1H NMR signals ranging from 5 to 3 ppm correspond to physisorbed water as well as to labile rapidly exchanging weakly hydrogen-bonded hydroxyl groups. The broad peak ranging from 8 to 2 ppm corresponds to silanol protons in a variety of hydrogenbonding environments (the lower the isotropic chemical shift, diso, the weaker the hydrogen bond) and the sharp peaks ranging from 1 to 2 ppm were assigned to non-hydrogen bonded silanol [47–60]. Moreover, in a recent study on precipitated silica nanoparticles [60] the 1H resonance at 1.1 ppm was tentatively assigned to inaccessible isolated silanol groups. The 1H MAS spectrum of pristine silica is shown in Figure 17.10. Since no annealing was undertaken to remove water, the contribution of the absorbed water at diso ¼ 4 ppm is the largest one. The occurrence of physically absorbed water molecules on the silica surface was shown by their removal, i.e. the weight loss that occurs as soon as the temperature is raised (30–100 C) (Figure 17.8). Owing to this large peak, the hydrogen-bonded silanols, which correspond to a distribution of chemical shifts between 2.5 and 8 ppm, cannot easily be seen. The NMR lines at 1.1 and 2.2 ppm are assigned to nonhydrogen-bonded silanols. The 1H-29Si CP-MAS NMR spectrum (Figure 17.11) of pristine silica exhibits the wellestablished three lines previously observed in amorphous silicas [11, 50–56, 60–67]. The peak at ca. –101 ppm is due to silicon with only one hydroxyl group, O3/2SiOH, often referred to as Q3 (Qn represents the SiO4 tetrahedron of the amorphous network, which forms n bonds with neighbouring tetrahedra). The resonance at ca. –111 ppm is assigned to siloxanes, O4/2Si (Q4), whose 29Si nuclei are crosspolarized by nearby protons leading to a resonance on the 1H-29Si CP-MAS NMR spectrum. The peak at ca. –92 ppm is assigned to silicon atoms having two geminated hydroxyl groups attached: O2/2Si(OH)2 (Q2). The isotropic chemical shifts, the line widths, and the relative intensities of these resonances are gathered in Table 17.5. Whereas IR spectroscopy cannot differentiate isolated and geminated hydroxyl groups [28], the occurrence of these latter species is clearly evidenced on the 1H-29Si CPMAS NMR spectra. On the other hand, IR spectroscopy allows differentiating isolated, terminal, and internal Si-OH groups. These NMR experiments confirm the IR spectroscopy observations concerning the heavily hydroxylated surface of this pristine silica: the studied sample presents more Q3 and Q2 species than the silica studied by Hartmeyer et al. [15, 60] or Barabash et al. [11]. This may partially explain the high fluorine ratios obtained after fluorination, since a proposed fluorination mechanism consists of a substitution of F ions for OH groups [28]. 17.6.2

Effect of Fluorination on the Nuclei Environments

In Figure 17.10 [71] the reduction in the signal-to-noise ratio may be related to the decrease of the amount of protons when the F content increases from pristine silica to 50F2 sample.

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535

Moreover, the intensity of the peak at 4 ppm, corresponding to the absorbed water, clearly decreases. As previously shown by IR spectroscopy [28], the amount of water rapidly decreases upon fluorination leading to the conclusion that fluorination improves the hydrophobic character of silica. The 1H MAS NMR spectra also show the decrease of the amount of both the hydrogenbonded and non-hydrogen-bonded silanols as the amount of grafted fluorine increases. IR spectroscopy, which allows the differentiation of isolated silanol groups from internal (nonaccessible) silanol groups, revealed internal Si-OH groups unaffected by F2 treatment (3660 cm1 band) whereas the intensity of the (OH) band of the isolated Si-OH group becomes weaker upon fluorination and finally vanishes for the highest F contents [28]. From the 1H MAS NMR spectra, the F/OH substitution that occurs during the fluorination process affects both the hydrogen-bonded (from 2.5 to 8 ppm) and nonhydrogen-bonded silanols (at 1.1 and 2.2 ppm). This result is surprising as the 1H resonance at 1.1 ppm was assigned to nonaccessible isolated silanol groups in a recent study on precipitated silica nanoparticles [15]. Finally, in agreement with the FTIR study and the F-titration data (Table 17.1), silanol species (hydrogen-bonded and non-hydrogen-bonded silanols as shown by NMR, and hydrogen-bonded and internal silanols as evidenced by FTIR) remain even for the highest F contents. H2O

Hydrogen bonded silanols

Non-hydrogen bonded silanols

50F2

10F2 SiO2 15

10

5

0

–5

δiso (ppm)

Figure 17.10 1H MAS (15 kHz) Hahn-echo NMR spectra of pristine (SiO2 sample) and fluorinated silica (10F2 and 50F2 samples). The double arrows indicate the chemical shift ranges of non-hydrogen-bonded silanols, physisorbed water and hydrogen bonded silanols. Reproduced by permission J. Phys. Chem. C 2009, 113, pp. 1865218660, American Chemical Society

Figure 17.11 displays the 1H-29Si CP-MAS NMR experimental and reconstructed spectra of fluorinated silica. It may be outlined that the number of transients was increased with the fluorination level in order to keep constant the signal-to-noise ratio. This mirrors the large decrease of the intensity showing the decrease of the silanol content from SiO2 to 50F2 samples in good agreement with the 1H NMR results. Even at the lowest fluorine

536

Functionalized Inorganic Fluorides

content, the Q2 geminated hydroxyl groups (ca. –92 ppm) disappear. On the contrary, the Q3 species (ca. –101 ppm) are still present but their relative intensity decreases when the F content increases. These data indicate the lower stability of the Q2 species upon fluorination. According to Hartmeyer et al. [60] who studied silica fluorinated by aqueous NH4F solution, the identified three lines (Figure 17.11 and Table 17.5) correspond to O2/2SiF2, O3/2SiF species, and pentacoordinated O4/2SiF groups, respectively. These assignments are deduced from a detailed NMR investigation, and supported by previous studies about fluorine-doped silica glasses [65, 66], silica fluorinated by a nonaqueous solution of NH4F [11] and fluorinated zeolites [68, 69].

50F2 exp. cal.

exp.

SiO2

cal.

–70 –80 –90 –100 –110 –120 –130 –140 –150 δiso (ppm)

Figure 17.11 Experimental and calculated 1H-29Si CP-MAS (5 kHz) NMR spectra of pristine (SiO2 sample) and fluorinated silica (50F2) (individual contributions are shown). Reproduced by permission J. Phys. Chem. C 2009, 113, pp. 1865218660, American Chemical Society

For highest fluorine contents (50F2), the reconstructions need two supplementary lines at ca. –124 and ca. –132 ppm (Figure 17.11 and Table 17.5). These two fluorosilicate species are related to the high fluorine content of the studied samples. The 19F-29Si CP-MAS NMR spectra of fluorinated silica, recorded with different contact times (Figures 17.12: contact time ¼ 5 ms) confirm the occurrence of five kinds of fluorosilicate species, and especially the two unprecedented species corresponding to the lines at ca. –124 and ca. –132 ppm. The relative intensities of these five lines increase from the 1H-29Si to the 19F-29Si CP-MAS NMR spectra confirming their attribution to fluorosilicate species. The same conclusion may be drawn from the comparison between

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Table 17.5 Isotopic chemical shifts (diso, ppm), line widths (L.W., ppm) and relative intensities (Int., %) as deduced from the reconstruction of the 1H-29Si CP-MAS NMR spectra of pristine and fluorinated silica and line assignments 50F2

SiO2 diso 92.0

L. W.

Int.

7.5

10

101.0

6.0

50

111.0

8.0

40

Assignment

diso

L. W.

Int.

96.5 101.5 107.0 112.0 118.5 124.5

8.0 7.0 6.5 7.5 7.0 4.5

9 22 8 54 5 1

O2/2Si(OH)2 O2/2SiF2 O3/2SiOH O3/2SiF O4/2Si O4/2SiF O3/2SiF2

the two 19F-29Si CP-MAS NMR experiments at different contact times: the shorter the contact time, the higher the fluorosilicate species relative intensity.

exp.

50F2

cal. –70 –80 –90 –100 –110 –120 –130 –140 –150 δiso (ppm)

Figure 17.12 Experimental and calculated 19F-29Si CP-MAS (5 kHz) NMR spectra of a fluorinated silica (50F2) recorded with a contact time of 5 ms (individual contributions are shown). Reproduced by permission J. Phys. Chem. C 2009, 113, pp. 1865218660, American Chemical Society

The occurrence of these new species can be related to the high fluorine level reached in these samples, compared to those measured by Hartmeyer et al. (from 0.5 to 14.6 wt % but with a large amount of (NH4)2SiF6 for the high fluorine levels)[60] or Barabash et al. (from the low signal-to-noise ratio of the recorded spectra) [11]. What could be the formula for these two new species? Four species are possible a priori: one tetracoordinated O1/2SiF3 species (mentioned by Barabash et al. [11]) and three pentacoordinated species O3/2SiF2, O2/2SiF3, and O1/2SiF4. The former O1/2SiF3 and the latter O1/2SiF4 present a low stability and would easily lead to SiF4 or SiF62. Moreover, it seems difficult to conceive the occurrence of O1/2SiF4 without forming O3/2SiF2 and O2/2SiF3 species. The occurrence of hydroxyfluorosilicate species, such as O3/2SiOHF or

538

Functionalized Inorganic Fluorides

O2/2SiOHF2 species for instance, is unlikely because of the strong probability of forming HF with two adjacent F and OH ligands. Then the most probable formula for these new species are O3/2SiF2 and O2/2SiF3.

50F2

exp. cal.

–120

–130

–140 –150 –160 δiso (ppm)

–170

–180

Figure 17.13 Experimental and calculated 19F MAS (25 kHz) NMR spectra of a fluorinated silica (50F2) (the individual contributions are shown). Reproduced by permission J. Phys. Chem. C 2009, 113, pp. 1865218660, American Chemical Society Table 17.6 Isotopic chemical shifts (diso, ppm), line widths (L.W., ppm) and relative Intensities (Int., %) of the NMR resonances as deduced from the reconstruction of the 19F MAS NMR spectra of 50F2 fluorinated silica; line assignment and relative amount (mol %) of the fluorosilicate species 50F2 diso 160.0 156.0 153.0 148.5 143.5 136.5 128.5 a

L. W.

Int.

Assignment

3.0 5.0 5.1 7.2 7.3 8.0 4.6

4.5 19.1 27.6 30.5 16.1 1.7 0.5

O2/2SiF2 O3/2SiFa O4/2SiF O3/2SiFb O3/2SiF2 O2/2SiF3 SiF6

mol % 2.6 21.6 31.3 34.6 9.1 0.7 0.1

isolated O3/2SiF species, bO3/2SiF species close to other groups of the same type.

Figure 17.13 displays the 19F MAS NMR spectrum of 50F2 fluorinated silica reconstructed with seven lines (Table 17.6). The method used for 1H-29Si and 19F-29Si was also applied, i.e., starting the reconstruction with the well-known and previously assigned 19F NMR lines. The resonance at diso(19F) ¼ 127 ppm can be assigned to SiF62 species [5, 60, 11] and its relative intensity indicates a very low amount, ca. 0.1 to 0.2 mol % (Table 17.6). This explains why this species could not be detected on the 19F-29Si CP-MAS NMR spectra nor by IR spectroscopy [28]. The three resonances at diso(19F)¼ 156, 153, and 149 ppm were previously observed in fluorinated silica and fluorinated zeolites.

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According to Hartmeyer et al., these resonances are assigned to isolated O3/2SiF groups, pentacoordinated O4/2SiF groups, and O3/2SiF species close to other groups of the same type, respectively (Table 17.6). On the other hand, no 19F NMR line was assigned to the O2/2SiF2 species. From the 1H29Si and 19F29Si CPMAS spectra, this species appears at low fluorine content simultaneously with the disappearance of the O2/2Si(OH)2 species and its relative amount remains roughly unchanged. On the 19F NMR spectra, the resonance at –160 ppm offers the same trend. Moreover, ab initio 19F chemical shift calculations performed by Liu et al. [70] to investigate the nature of possible species in aluminosilicate glasses lead to a diso(19F) value of 155 ppm for O2/2SiF2. Therefore the line at diso(19F) ¼ 160 ppm should be assigned to O2/2SiF2. The two remaining lines at diso(19F) ¼ 143.5 and 136.5 ppm should then correspond to the two unprecedented species, O3/2SiF2 and O2/2SiF3. Given that the less fluorinated species presents the highest abundance, the higher relative intensities of the 29Si lines at ca. 124 ppm and of the 19F lines at 143.5 ppm allow assigning these lines to the O3/2SiF2 species. Consequently, the 29Si line at ca. 132 ppm and the 19F line at 136.5 ppm are assigned to the O2/2SiF3 species. These assignments are supported by the ab initio 19F chemical shift calculations of Liu et al. which give diso values ranging between 128 and 132 ppm for O2/2SiF3. The 19F NMR spectra being quantitative, the relative amount of each species has been calculated (Table 17.6). As expected, the relative amount of O3/2SiF species as a close neighbour of other O3/2SiF groups increases whereas the relative amount of isolated O3/2SiF species decreases when the F2 content increases. Actually, the relative amount of O2/2SiF2 species which correspond formally to neutral entities remains quasi-constant whatever the F2 content in the gas mixture whereas the relative amount of isolated O3/2SiF and pentahedral O4/2SiF species globally decreases and the relative amount of pentahedral O3/2SiF2 and O2/2SiF3 species as well as O3/2SiF species close to other groups of the same type increases with the F2 concentration. The F2-direct fluorination with higher F2 contents allows the concentration of isolated monofluorinated species to be limited in favour of di- or trifluorinated species, as well as interactions between O3/2SiF species. However, it seems surprising to stabilize more pentahedral O3/2SiF2 species than O2/2SiF2 species. This means that multi Si-F bonds up to 3 can be stabilized around each Si atom when the number of siloxane bridges is equal to 2 or 3. Moreover no O1/2SiF3 species have been detected. The relative amount of 4/2SiF fluorinated species decreases versus the F2 gas concentration whereas the content of O3/2SiF entities increases with the F2 gas rate. Then, in addition to the redox mechanism involving F2/F and O2/OH couples, the high oxidizing power of F2-gas allows the breaking of one Si-O-Si bridge replaced by an additional Si-F bond. The occurrence of pentahedral species such as O2/2SiF3, O3/2SiF2, or O4/2SiF can be easily explained by taking into account the preferred stabilization of [SiO4/2] tetrahedra and [SiF6]2octahedra in oxides and fluorides, respectively. Finally for the following entities [O2/2SiF3], [O3/2SiF2], and [O4/2SiF], the effective charge is always equal to –1, intermediate between the neutral [SiO4/2] tetrahedron and the [SiF6]2 octahedron. The polarization of such pentahedral species should strongly increase with the number of fluorine ligands due to the coexistence of various species with different electronegativities. This explains the decrease of the thermal stability for highly fluorinated silica leading to the formation of gaseous SiF4, as well as the increase of the water affinity with the F2 content in the gaseous mixture due to the high concentration of highly polar species.

540

17.7

Functionalized Inorganic Fluorides

Conclusions on the F2-gas Fluorination Mechanism of Mesoporous Silica

At present, direct F2 fluorination is the only route to access to highly fluorinated silica, i.e. up to 13 wt % [28]. Fluorine is substituted for hydroxyl groups at the surface of mesoporous silica by suitable gaseous treatments using elemental F2. Pre-treatments of starting silica, outgassing or annealing at higher temperature under Ar, which lead either to remove water (T ¼ 200 C) or to largely decrease the hydroxyl content (T ¼ 400 C600 C), allow to tune the amount of Si-F grafting during the fluorination process. The other key parameters of this synthesis route are the fluorination temperature (25 C and 100 C) and the F2/Ar rate in the gas mixture. However, whatever the synthesis conditions, an amount of 0.11 OH groups per Si atom, nonaccessible to elemental F2 and corresponding to internal hydroxyls mainly located into the silica framework is systematically present whatever the fluorination treatment. The SiO2xy(OH)2xF2y, chemical composition, identified on the basis of F titration and FTIR analysis varies from SiO1.75OH0.50 starting material up to SiO1.74(OH)0.09F0.43 composition. Despite the reduction of surface area after the fluorination treatment, the mesoporous nature of the material remains almost identical. The F content is higher at the surface of the agglomerates of the mesoporous silica than in the bulk and this difference is governed by the surface area showing that F2 directfluorination proceeds mainly at the surface. However for specific pre-treatment temperatures lower than T ¼ 600 C as well as fluorination temperatures lower than T ¼ 100 C, bulk and surface F/Si atomic ratios tend almost to the same value, demonstrating that the F2-direct route leads also to rather homogeneous materials. For instance, if the fluorination treatment occurs at T ¼ 100 C, or if silica is pre-treated at T ¼ 200 C followed by a direct fluorination at RT, the Si-F species are equally distributed between surface and bulk. Actually the limitation of physisorbed and structural water trapped into mesoporous silica, as well as the occurrence of large amount of isolated, terminal and bound OH groups allow increasing the number of Si-F chemical bonds. Taking into account a redox process involving O2/OH and F2/F couples, the Le Chatelier law perfectly accounts for the key role of H2O and OH, and for the large content of stable Si-F bonds obtained through this process. This redox mechanism proceed up to 50 % F2/Ar gas mixture; for higher F2gas contents in the gaseous mixture, the SiO2 etching takes place with a partial breaking of siloxane bridges and becomes dominant. The two reaction paths probably compete but in both these mechanisms the formation of HF cannot be excluded. From a thermodynamic point of view, one can assume that the first (low F2 concentration) reaction associated to a redox process involves the two half redox equilibria with F2, O2 and H2O as gas phases, as well as silanol and Si-F chemical bondings in solid state phases [28]. The elimination of physisorbed water trapped into silica and the occurrence of high hydroxyls content and F2 concentration contribute to a right shift of the first reaction and the creation of Si-F chemical bonds. Thus, a fluorination treatment at room temperature with 50 % F2/Ar of silica outgassed at T ¼ 200 C, or the fluorination at T ¼ 100 C with 20 % F2/Ar of not-outgassed as-received silica, yields a maximum content of grafted fluorine. This first process could be thus associated to an exothermic redox reaction involving F2 gas and O2 in addition with H2O gas evolution. This result differs from what generally observed with aqueous organic fluorination routes which imply in a

Access to Highly Fluorinated Silica

541

simple manner SiO2 and H2SiF6 as well as Brønsted acid/base equilibrium. Considering this redox process, it is then clear why the elemental F2 fluorination allows controlling the F/Si atomic ratio by decreasing the water content or increasing the OH concentration, the OH groups being systematically removed, leading probably to the O2 evolution. However it has been pointed out in the experimental part that under fluorinated atmospheres, the decomposition of silica into SiF4 could not be totally ignored. In particular it could be assumed that in highly reactive conditions (high F2 concentration), the formation of volatile SiF4 occurs in more noticeable amounts. The subsequent volatilisation of this gaseous species breaks numerous siloxane bridges thus opening new paths through to the core of the agglomerates. The created anionic vacancies are immediately filled with fluorine, which is present in large amounts in the reactive atmosphere, giving rise to very high levels of bulk fluorination. Various [SiF6]2 octahedral and [O(4n)/2SiFn] (1£ n £3) tetrahedral species or even fivefold coordinated [O4/2SiF] environments have been identified by NMR spectroscopy (1H,19F,29Si) and FTIR spectroscopy [11, 15, 25]. Besides various fluorosilicate species previously observed in moderately fluorinated silica such as O2/2SiF2, O3/2SiF and O4/2SiF, two unprecedented species have been evidenced for highly fluorinating conditions [71]. The 19F (diso19F ¼ 143.5 ppm and 136.5 ppm) and 29Si (diso29Si ¼ 124 ppm and 132 ppm) NMR resonances of these two species have been tentatively assigned to O3/2SiF2 and O2/2SiF3 based on the relative intensities of 19F and 19F29Si NMR lines and previous ab initio 19F isotropic chemical shift calculations. The physico-chemical properties of silica powder can be thus modified by fluorine substitution. The F2-gas direct fluorination of silica leads to a substitution of Si-OH groups by Si-F and SiF2 species, and to unprecedented pentahedral species O2/2SiF3 and O3/2SiF2 by breaking various siloxane bridges. The high oxidizing power of F2 gas allows such unusual species to be reached due to a coupling between a redox reaction and an etching phenomenon. Moreover, fluorination conditions lead to fluorinated products stable up to 400 C. In addition, such fluorinated silica exhibit a lower water uptake than untreated silica in highly humid environment (75 % relative humidity) for long duration times (up to 50 days), a property that can be correlated to its hydrophobic character. It can be established that a reduction of concentration of silanol groups on the silica surface is generally accompanied by an increase of the hydrophobic character. In molecular adsorption on a solid surface, the chemical nature of the solid surface plays an important role and hydrophilic-hydrophobic properties are indeed decisive in the adsorption control of polar/ non-polar molecules.

Acknowledgements We gratefully acknowledge the European Community for the financial support in the STREP FUNFLUOS (FUNctionalized FLUOrideS) network (NMP3-CT-2004-505575) The IR investigations were carried out in collaboration with A. Vimont and M. Daturi at Univ. Caen, France and NMR results were obtained by C. Legein and J.Y. Buzare´ at the University of Le Mans, France. C. Labrugere is acknowledged for XPS experiments and fitting.

542

Functionalized Inorganic Fluorides

References [1] Barly, D. Silicas.Characterization of Powder Silicas.Academic Press,New York, 1976. [2] Iler, R.K.The Chemistry of Silica.Wiley-Interscience,New York, 1979. [3] Bossaert, W.D.;De Vosq, D.E.;Van Rhijn, W. M.;Bullen, J.;.Grobet, P.J.;Jacobs, P.A.J. Catal. 1999, 182, 156. [4] Chapman, I. D.;Hair, M. L.J. Catal. 1963, 2,145. [5] Duke, C.V.A.;Miller, J. M.;Clark, J. H.;Kybett, A. P.Spectrochim. Acta, Part A 1990, 46,1381. [6] Branda, M. M.;Montani, R.A.;Castellani, N.J. Surf. Sci. 2000, 446 L89. [7] Vansant, E. F.;Van Der Voort, P.;Vrancken, K. C.Characterization and Chemical Modification of the Silica Surface, Elsevier, Amsterdam, 1997 (Chapters 1–6). [8] Schwertfeger, F.; Emmerling, A.; Gross, J.; Schubert, U.; Fricke, J. Y. A. Attia (Ed),Sol-gel Processing and Applications, Plenum, New-York, 1994, 343. [9] Katsuo, T.;Satohito, Y.;Kimio, T.Bull. Jpn. Pet. Inst. 1970, 12, 136. [10] Pesek, J. J.; Matyska, M. T.; Abuelafiya, R. R. Chemically Modified Surfaces: Recent Developments, Royal Society of Chemistry, Cambridge, UK, 1996. [11] Barabash, R. M.;Zaitsev, V. N.;Kovalchuk, T. V.;Sfihi, H.;Fraissard, J.J. Phys. Chem. A, 2003, 107, 4497–4505. [12] Campostrini, R.;Ischia, M.;Carturan, G.;Armelao, L.J. Sol-Gel Sci. and Tech. 2002, 23, 107–117. [13] Li, W.; Willey, R. J. J. Non-Cryst. Sol. 1997, 212, 243–249. [14] Lebeau, B.; Marichal, C.; Mirjol, A.; Soler-Illia, G. J. de A. A.; Buestrich, R.; Popall, M.; Mazerolles, L.; Sanchez, C. New J. Chem. 2003, 27, 166–171. [15] Hartmeyer, G.; Marichal, C; Lebeau, B.; Rigolet, S.; Caullet, P.; Hernandez, J. J. Phys. Chem. C. 2007, 111, 9066–9071. [16] Kang, J. K.; Musgrave, C. B. J. Chem. Phys. 2002, 116, 275–280. [17] Miyajima, H.; Katsumata, R.; Nakasaki Y.; Nishiyama Y.; Hayasaka N., J. Applied Physics 1996, 35 (12A), 6217. [18] Haung, Y; Sarker, A.; Schultz, P. J. Non-Cryst. Solids 1978, 27, 29. [19] Kirchhof, J.; Unger, S.; Klein, K. F.; Knappe, B. J. Non-Cryst Solids 1995, 181, 266. [20] Tsukuma, K.; Yamada, N.; Kondo, S.; Honda, K.; Segawa, H. J. Non-Cryst Solids 1991, 127, 191. [21] Coburn, W.; Chen, M. J. Appl. Phys. 1980, 51, 3134. [22] Donnelly, M.; Flamm, D. L. J. Appl. Phys. 1980, 51, 5273. [23] Yoshimaru, M.; Koizumi, S.; Shimokawa, K. J. Vac. Sci. Technol. A 1997, 15 (6), 2915. [24] Yoshimaru, M.; Koizumi, S.; Shimokawa, K. J. Vac. Sci. Technol. A 1999, 17 (2), 425. [25] Youngman, R. E.; Sen, S. J. Non-Cryst Solids 2004, 349, 10–15. [26] Paul, M. C.; Sen, R.; Bandyopadhyay, T. J. Mater. Sci. 1997, 32, 3511. [27] Yoo, W. S.; Swope, R.; Sparks, B.; Mordo, D. J. Mater. Res.1997, 12 (1), 70. [28] Lataste, E.; Demourgues, A.; Leclerc, H.; Goupil, J.-M.; Vimont, A.; Durand, E.; Labruge`re, C.; Benalla H.; Tressaud A. J. Phys. Chem. C 2008, 112, 10943. [29] Grannec, J.; Lozano, L. Preparative Methods in Inorganic Solid Fluorides, Ed: P. Hagenmuller, Academic Press. New York, USA, 1985, 18. [30] Hartmann, S. R.; Hahn, E. L. Phys. Rev. 1962, 128, 2042. [31] Pines, A.; Gibby, M. G.; Waugh, J. S. J. Chem. Phys. 1972, 56, 1776. [32] Seel. F; Angew. Chem. Internat edit. 1953, 3(6) 424. [33] Paparazzo, E. Surface and Interface Analysis 1996, 24, 729–730. [34] Barrett E.P., Joyner L.G.; Halenda P.H. J. Amer. Chem. Soc. 1951, 73, 373. [35] Sindorf, D. W.; Maciel, E. J. Phys. Chem. 1983, 87, 5516–5521. [36] Gallas, J. P.; Lavalley, J. C.; Burneau, A.; Barres, O. Langmuir 1991, 7, 1235–1240. [37] Li, W.; Willey, R. J. J. Non-Cryst. Solids 1997, 212, 243–249. [38] Carteret, C. Ph.D. Thesis: ‘Etude par spectroscopie dans le proche infra-rouge et mode´lisation des structures de surface et de l’hydratation de silices amorphes’, 1999, University of Nancy I, France. [39] Ek S.; Root A.; Peussa M.; Niinistø L. Thermodynamica Acta 2001, 379, 201–212. [40] Hirschfeld T. and Fateley W.G. Applied Spectroscopy 1977, 31 (1), 42–43.

Access to Highly Fluorinated Silica [41] [42] [43] [44] [45] [46] [47] [48] [49] [50] [51] [52] [53] [54] [55] [56] [57] [58] [59] [60] [61] [62] [63] [64] [65] [66] [67] [68] [69] [70] [71]

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18 Preparation and Properties of Rare-earth-Containing Oxide Fluoride Glasses Susumu Yonezawa, Jae-ho Kim and Masayuki Takashima Graduate School of Engineering, University of Fukui, Bunkyo 3–9–1, Fukui 910–8507, Japan

18.1

Introduction

Oxide fluoride glasses contain two different anions that have different valence electrons and different degrees of polarization. It is interesting to compare properties of the oxide fluoride glass with those of the oxide or fluoride glass. No reports have described a glass that can contain a rare-earth fluoride with such high contents. The properties of oxide fluoride glasses have not yet been summarized systematically, as they have been for oxide glasses or fluoride glasses. It is interesting to study preparation processes and characteristics of these glasses to develop new functional materials. We have reported the preparation and properties of oxide fluoride glasses containing rare-earth elements [1–4]. The oxide fluoride glasses are anticipated as a new optical material. Every rare-earth element has unique optical properties because of its arrangement of electrons in the 4f orbital. It is important to find matrices into which these rare earth elements can be doped and that can exhibit their performance to develop new optical and magnetic materials. During the preparation process [5] of some materials of solid electrolyte containing rare earth elements, one type of glass was obtained by chance. Based on quantitative analyses using EPMA, this glass contained all the starting elements (O, F, and two species of rare

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

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earth elements). In addition, Si and Al, which were not contained in starting materials, were detected in the glass. The binary rare-earth oxide fluoride solid electrode was prepared using a solid state reaction between a rare earth oxide (Ln2O3) and a rare earth fluoride (Ln’F3) at a temperature higher than 1000 C. During the mixing process, an agate ball mill was used to mix the starting materials. During calcination, the mixture was heated to over 1000 C after it was packed into an aluminium tube. The Si and Al in the glass had to be supplied from these materials as glass network formers. This glass contained more than 70 wt % of rare earth elements, as inferred from the results of the analysis using EPMA (fundamental parameter method (FP method)). Actually, the FP method is the quasi-quantitative analysis method using the instrument’s standard library for the peak intensities of the elements. The resultant glass was a new oxide fluoride glass containing large amounts of rare-earth elements. Several reports have described preparation of glasses containing rare earth elements for use as optical or magneto-optical materials [6–10].

18.2

Preparation and Basic Characteristics of Oxide Fluoride Glasses Containing LnF3

This section presents a description of the preparation process and basic characterization of oxide fluoride glasses containing CeF3 and NdF3 (light rare earth), TbF3 (middle rare earth), and HoF3 (heavy rare earth). Based on those results, preparation methods were extended to other rare-earth elements in the lanthanide series. Properties of those glasses were compared and summarized. 18.2.1

Preparation of Oxide Fluoride Glasses Containing LnF3

The glass in the NdF3-SiO2-Al2O3 system has been obtained once. However, the reproducibility of the synthesis was not confirmed. Detailed analyses of the glass revealed that AlF3 was contained in the product in the NdF3-SiO2-Al2O3 system, meaning that some part of NdF3 was hydrolysed to form HF; this HF has reacted with Al2O3 to form AlF3. Controlling the content of AlF3 is apparently an important factor to prepare the oxide fluoride glass reproducibly. However, the higher temperature for melting the sample causes large variation of the AlF3 content in the product because AlF3 sublimes at a temperature that is remarkably higher than 700 C. In addition, rare-earth trifluorides readily undergo pyrohydrolysis to form their oxides at temperatures higher than 1000 C [11]. The network forming oxide GeO2 was chosen because it has the lowest melting point among SiO2 (1730 C), Al2O3 (2045 C), and GeO2 (938 C) in this study. In addition, various fluorides were tested as glass network modifiers and additives to lower the melting point. Consequently, the oxide fluoride glasses containing light rare earth (NdF3) were obtainable reproducibly in the system of NdF3-AlF3-GeO2. The glass was analysed using the ZAF method with EPMA, which showed clearly that the obtained glass was incapable of maintaining its nominal composition. The Al content, especially, tended to decrease from the starting mixture ratio. The Al contents in the glass never became greater than 34 mol %, even if the starting material contained more than 34 mol % Al. The Nd content

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reached 63 mol % as the maximum value. Glasses with a rare earth content higher than 30 mol % have not been reported to date. These oxide fluoride glasses with a high content of rare earths are anticipated as new functional materials. Figure 18.1 presents a phase diagram of TbF3-AlF3-GeO2. The maximum TbF3 content was 50 mol % in the glass. Reportedly, rare earth trifluorides hydrolysed at a temperature higher than 700 C [11,12]; a lower melting temperature is preferred in order to avoid composition changes during high-temperature processing. Figure 18.2 shows that the results of quantitative analyses obtained by EPMA measurement of TbF3-AlF3-GeO2 glasses indicated that the cationic compositions in the glass deviated from that of the starting ratio. Furthermore, the composition change in the glass increased concomitantly with increasing AlF3 content. This system contained no glass with 50 mol % TbF3, even though the glass was obtained from a starting mixture containing 50 mol % TbF3. The amount of AlF3 must be optimized to realize good reproducibility because the excess of AlF3 drastically alters the glass composition from the starting ratio. The BaF2 was added to starting materials instead of AlF3 because BaF2 sublimates only slightly, but hydrolyses at melting temperatures of 1200 C. Figure 18.3 presents a phase diagram of the NdF3-BaF2-GeO2 system. This system had a large glass formation composition area. The NdF3 content in the glass reached 40 mol % as a nominal composition. Figure 18.4 presents the cationic composition in the glasses of NdF3-BaF2-GeO2 system resulting from quantitative analyses using EPMA. The glass obtained in this system maintained the starting cationic compositions better than the NdF3-AlF3-GeO2 system did. Apparently, by adding BaF2, the oxide fluoride glasses were obtained reproducibly and the nominal composition was maintained in the glass product. Overall, BaF2 was a better component to prepare the oxide fluoride glass containing NdF3 than AlF3 was. TbF3 / mol%

AlF3 / mol%

GeO2 / mol%

Figure 18.1 Phase diagram of TbF3-AlF3-GeO2 system shown by nominal composition. Closed and open circles respectively represent crystal and glass phases. Reproduced from [2] by permission of Elsevier

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Functionalized Inorganic Fluorides Tb / mol%

Al / mol%

Ge / mol%

Figure 18.2 Cationic composition in the glass of TbF3-AlF3-GeO2 system glasses measured using EPMA. Reproduced from [2] by permission of Elsevier NdF3 / mol%

BaF2 / mol%

GeO2 / mol%

Figure 18.3 Phase diagram of NdF3-BaF2-GeO2 system shown by nominal composition. Closed and open circles respectively represent crystal and glass phases. Reproduced from [2] by permission of Elsevier

Based on results of the glass preparation in the ternary system glass, preparation of 40LnF3-20BaF2-40GeO2 (Ln; La–Nd, Sm–Lu) glasses by melting at 1200 C for 1.5 h was attempted. Only in cases of NdF3, SmF3, EuF3, and GdF3 (middle rare-earth fluorides) were glasses obtained.

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Nd / mol%

Ba / mol%

Ge / mol%

Figure 18.4 Cationic composition in the glass of NdF3-BaF2-GeO2 system measured using EPMA. Reproduced from [2] by permission of Elsevier

Choosing the system containing HoF3, a glass-formation process with a heavy rare-earth trifluoride was investigated. Several chemicals were tested for use as additive components to obtain oxide fluoride glasses at lower melting temperatures and with lower weight loss during melting. Finally, results show that both BaF2 and AlF3 are necessary to obtain, reproducibly, the glasses containing HoF3; in addition, accurate controls of their contents are needed. Figure 18.5 portrays the glass-forming condition relative to the quantity of AlF3 and melting temperature. The oxide fluoride glasses were obtainable reproducibly by adding more than 6 mol % AlF3; excessive AlF3 changed the glass composition, as described above. The AlF3 content would be better maintained at ca. 10 mol % to obtain glass reproducibly without a composition change through the melting process. It was possible to prepare the glass at 1175 C in case of GeO2 system, although about 1300 C was needed in the case of the SiO2 system. Figure 18.6 portrays a phase diagram of the HoF3-BaF2-AlF3-GeO2 system in which the AlF3 content was fixed at 10 mol %. This diagram shows that this glass system has a wide composition range for glass formation. An oxide fluoride glass was prepared with maximum content of HoF3 of 50 mol %. Figure 18.7 presents results of quantitative analyses of HoF3-BaF2-AlF3-GeO2 glasses using EPMA (ZAF method). Consequently, only slight deviation of the cationic composition in the glass product from the nominal composition of the starting mixture was recognized. Regarding 10HoF3-10BaF210AlF3-70GeO2 and the 50HoF3-10BaF2-10AlF3-30GeO2 glasses, the respective analytical cationic ratios in the product glasses were Ho: Ba: Al: Ge ¼ 6.1: 9.0: 11.6: 73.2 and 41.4: 13.6: 8.8: 36.1. These values approximately agreed with the nominal composition. Finally, the glasses were obtained reproducibly in the LnF3-BaF2-10AlF3-GeO2 system for Ln ¼ La – Lu (except for Pm).

550

Functionalized Inorganic Fluorides

1200

Temperature / °C

1190

1180

1170

1160

1150

1140

0

5

10

15

20

25

x / mol%

Figure 18.5 Classification of the products for 10HoF3-(30-x)BaF2-xAlF3-60GeO2 systems at various melting temperatures. * and respectively signify glass and crystalline phases. Reproduced from [2] by permission of Elsevier

•

HoF3 / mol%

BaF2 + 10AlF3 / mol%

GeO2 / mol%

Figure 18.6 Phase diagram of HoF3-BaF2-10AlF3-GeO2 shown by nominal composition. The AlF3 content was fixed at 10 mol%. Closed and open circles respectively denote crystal and glass phases. Reproduced from [2] by permission of Elsevier

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Ho / mol%

Ba + 10Al / mol%

Ge / mol%

Figure 18.7 Cationic composition in the glass of HoF3-BaF2- AlF3-GeO2 system measured using EPMA. Reproduced from [2] by permission of Elsevier

Regarding the CeF3-BaF2-10AlF3-GeO2 and CeF3-BaF2-10AlF3-SiO2 systems, the colour of the glasses prepared in the manner described above (in an argon atmosphere) was brown. This might result from the mixed valency of Ce3þ and Ce4þ because of the decomposition of CeF3 during the glass preparation process. To prepare pale yellow glasses in the CeF3-BaF2-AlF3-SiO2 system, the melting process was conducted in a CO atmosphere. At the composition, around 10CeF3-20BaF2-10AlF3-60SiO2 or 20CeF3-10BaF210AlF3-60SiO2, the glass was reproducibly prepared by heating at 1300 C for 90 min. Optimizing the conditions such as heating temperature, holding time and heating rate, and controlling the hydrolysis of rare earth fluoride, the composition range to prepare the glass was extended slightly. Brown glass tended to be obtained in cases of higher CeF3 contents, even in CO. It was possible to produce a glass containing 40 mol % of CeF3 as a maximum. During the glass-preparation process, the sample weight decreased about 18 wt %. Fluorine must be lost completely if this weight loss occurred because of the hydrolysis only. However, it was confirmed by X-ray fluorescent spectroscopy (XFS) measurement that 43 and 22 % of fluorine in the starting mixture remained respectively in the products prepared in CO and Ar atmospheres. Therefore, the hydrolysis of CeF3 was controlled in a CO atmosphere although CeF3 was hydrolysed during both preparation processes in CO and Ar atmospheres. The hydrolysis of LnF3 proceeds to produce Ln2O3 via LnOxFy [11]. For cerium, the final product of hydrolysis is CeO2, which might be derived from Ce2O3 and/or CeOxFy: the hydrolysis of CeF3 includes several reaction processes. For the glasses examined here, the reaction from CeOxFy to CeO2 is apparently controlled in a CO atmosphere. For melting in an Ar atmosphere, the oxidative decomposition of CeOxFy (reaction with H2O and O2 to produce HF) proceeds and Ce4þ is generated in the sample (x value increases). The total reaction from CeF3 to CeO2 might be characterized by the following equation: 2CeF3 þ

552

Functionalized Inorganic Fluorides

3H2O þ 1/2O2 ¼ 2CeO2 þ 6HF. The O2 that is ubiquitous in the atmosphere might be an oxidizer in this case. Furthermore, CO might greatly reduce the partial pressure of O2 in the atmosphere by CO þ 1/2O2 ¼ CO2 equilibrium.

18.2.2

Density and Refractive Index

Densities and refractive indexes of NdF3-Al2O3-SiO2, NdF3-AlF3-GeO2, NdF3-BaF2GeO2, TbF3-BaF2-AlF3-GeO2, HoF3-BaF2-AlF3-GeO2, and 10LnF3-20BaF2- 10AlF360GeO2 system glasses are presented in Figure 18.8. In that figure, (–) and (þ) respectively correspond to data of typical oxide and fluoride glasses taken from the database, INTERGLAD [13]. The oxide glasses in this case include quartz glass [14], borosilicate glass (COVER 18-18; Iwaki Glass Co. Ltd.), PbO-WO3-P2O5-CdO-TiO2 glass [15], and various oxide glasses in the literature found in the INTERGLAD database [13]. In addition, ZBLAN [16–21] glass and others found in the database were chosen as fluoride glasses. Curves (A) and (B) shown in Figure 18.8 were derived respectively by fitting the equation (n ¼ C1exp (C2d) where C1 and C2 are constants) to data of oxide glasses and fluoride glasses. These empirical curves emphasize characteristics of the oxide and the fluoride glasses. Plots of the oxide fluoride glasses are between curves (A) and (B). Especially, TbF3-BaF2-10AlF3-GeO2, HoF3-BaF2-10AlF3-GeO2, and LnF3-BaF2-10AlF3-GeO2 (Ln; Y–Nd, Sm–Lu) glasses used in this study exhibited a constant refractive index irrespective of density. The relation between density and refractive index is generally given as the Lorentz–Lorenz equation [15]. n2 1 4pN V¼ ¼RL n2 þ 2 3

(18:1)

Therein, n, V, N, , and RL respectively represent the refractive index, molecular volume, Avogadro’s number, polarizability, and molecular refraction. The Gladstone–Dale equation is derived from Equation (18.1) when n is close to 1, as shown below [15]. n1¼

R 2pN 2pN ¼ ¼ d ¼ RG d v v M

(18:2)

In those equations, M, d, and RG respectively denote the molecular weight, density, and the Gladstone–Dale constant. Generally, this equation is used to express the relation between density and the refractive index of a glass. The RG value varies with glass composition. The RG of the fluoride, the oxide, and the oxide fluoride glasses were, respectively, 0.9 104 – 1.4 104, 1.8 104 – 2.9 104, and 1.0 104 – 1.8 104. Data of the NdF3-Al2O3-SiO2 system closely approximated the curve (A) corresponding to oxide glasses in Figure 18.8. The glass network of NdF3-Al2O3-SiO2 was suggested to be an oxide such as Al2O3 and SiO2. However, in the case of NdF3-AlF3-GeO2 or NdF3-BaF2GeO2, systems in which the glass matrix consists of both an oxide and a fluoride, the plots of the refractive index against the density deviated from curve (A) to curve (B),

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2.2

Refractive index / –

2.0

1.8

1.6

1.4

1.2

2

3

4 5 6 Density / × 103 kg m–3

7

8

Figure 18.8 Relation between density and the refractive index for various glasses. Reproduced from [2] by permission of Elsevier. : LnF3-BaF2-AlF3-GeO2 (Ln; Y–Nd, Sm–Lu) : oxide glass* þ : fluoride glass* * : TbF3-BaF2-AlF3-GeO2 ~ : NdF3-Al2O3-SiO2 & : HoF3-BaF2-AlF3-GeO2 & : TbF3-BaF2-AlF3-SiO2 n : NdF3-AlF3-GeO2 ^ : NdF3-BaF2-GeO2 (* : data from INTERGRAD[14])

•

corresponding to that of the fluoride glasses. For LnF3-BaF2-AlF3-GeO2 glasses, the plots were located between curves (A) and (B). Data of glasses that consisted of 70 mol % oxide and 30 mol % fluoride (10HoF3-10BaF2-10AlF3-70GeO2) closely approximated curve (A) for oxide glasses. Data of glasses consisting of 30 mol % oxide and 70 mol % fluoride (50HoF3-10BaF2-10AlF3-30GeO2) closely approximated curve (B) for the fluoride glass area. Consequently, the oxide fluoride glasses were located at the intermediate area in the relation between the refractive index and the density. No data for the simple oxide or fluoride glass was reported in this region previously.

18.2.3

Glass Transition Temperature

Respective glass transition temperatures of the HoF3-BaF2-10AlF3-GeO2 glasses, TbF3-BaF2-AlF3-GeO2 glasses, TbF3-BaF2-10AlF3-SiO2 glasses, and 10LnF320BaF2-10AlF3-60GeO2 glasses are presented in Tables 1–4. The glass transition

554

Functionalized Inorganic Fluorides

temperatures of BaF2-free glasses were higher than those of quaternary systems such as LnF3-BaF2-AlF3-GeO2 glasses. Glasses containing a larger amount of BaF2 as a network modifier tend to have lower glass transition temperatures because divalent ions such as Ba2þ cut the glass network. Fluoride glasses such as ZBLAN [16–21] exhibit a glass transition temperature around 300–400 C. Oxide fluoride glasses obtained in this study must be more thermally stable than the fluoride glasses. Table 18.1 Glass transition temperatures of HoF3-BaF2-AlF3-GeO2 glasses. Reproduced from [2] by permission of Elsevier HoF3:BaF2:AlF3:GeO2

Tg / oC

10:10:10:70 20:10:10:60 30:10:10:50 40:10:10:40 50:10:10:30 10:20:10:60 20:20:10:50 30:20:10:40 40:20:10:30 50:20:10:20 10:30:10:50 20:30:10:40

592.0 599.8 614.4 608.3 572.3 599.9 600.9 585.7 564.5 541.8 576.8 561.9

Table 18.2 Glass transition temperatures of TbF3-BaF2-AlF3-GeO2 glasses. Reproduced from [2] by permission of Elsevier TbF3:BaF2:AlF3:GeO2

Tg / oC

10:10:10:70 20:10:10:60 30:10:10:50 40:10:10:40 50:10:10:30 10:20:10:60 20:20:10:50 30:20:10:40 40:20:10:30 10:30:10:50 20:30:10:40 30:30:10:30 30:0:10:60 40:0:10:50 50:0:10:40 30:0:20:50 40:0:20:40 30:0:30:40 40:0:30:30

606.7 631.1 619.9 628.8 618.2 578.4 603.4 593.4 574.0 593.0 583.4 571.9 675.2 669.1 639.9 664 610.8 605.1 622.2

Rare-earth-Containing Oxide Fluoride Glasses

555

Table 18.3 Glass transition temperatures of TbF3-BaF2-AlF3SiO2 glasses. Reproduced from [2] by permission of Elsevier TbF3-BaF2-AlF3-SiO2

Tg / C

30-10-10-50 40-10-10-40 50-10-10-30 10-20-10-60 20-20-10-60 30-20-10-40

666.5 633.1 590.0 657.6 608.3 588.7

Table 18.4 Glass transition temperatures of LnF3-BaF2-AlF3-GeO2 (Ln; Y-Nd, Sm-Lu) glasses. Reproduced from [2] by permission of Elsevier

18.3

18.3.1

Composition

Tg / oC

10YF3-20BaF2-10AlF3-60GeO2 10LaF3 -10BaF2-10AlF3 -70GeO2 10CeF3-20BaF2-10AlF3-60GeO2 10PrF3-20BaF2-10AlF3-60GeO2 10NdF3-20BaF2-10AlF3-60GeO2 10SmF3-20BaF2-10AlF3-60GeO2 10EuF3-20BaF2-10AlF3-60GeO2 10GdF3-20BaF2-10AlF3-60GeO2 10TbF3-20BaF2-10AlF3-60GeO2 10DyF3-20BaF2-10AlF3-60GeO2 10HoF3-20BaF2-10AlF3-60GeO2 10ErF3-20BaF2-10AlF3-60GeO2 10TmF3-20BaF2-10AlF3-60GeO2 10YbF3-20BaF2-10AlF3-60GeO2 10LuF3-20BaF2-10AlF3-60GeO2

608.4 630.3 585.3 582.2 590.1 589.6 586.7 592.8 578.4 599.3 599.9 609.8 579.5 583.4 581.0

Optical and Magnetic Properties of LnF3-BaF2-AlF3-GeO2 (SiO2) Glasses Optical Properties of HoF3-BaF2-AlF3-GeO2 Glasses

Some reports have described studies of oxide fluoride, oxyfluoride, and fluorophosphate glasses [22–28]. Oxide fluoride glasses have been researched as host materials for optically active ions because they have low phonon energies that correspond to oxide glasses, and high chemical and mechanical stabilities related to fluoride glasses. Although oxide and fluoride ions have similar ionic radii, the ratio of oxide and fluoride ions in the glass must alter the coordination structure that affects the elements’ functionality because of their different valences. For example, binary rare earth metal oxide fluorides such as Nd2Eu2O3F6 have an ordered ionic configuration that engenders higher electric

556

Functionalized Inorganic Fluorides

conductivity than that of YSZ-11 [29]. Oxide fluoride materials containing multication species can exhibit unique properties because of their ordered-disordered ionic configuration. In this section, optical properties related to contents of LnF3 were reviewed using HoF3 as a probe. The results of absorption spectra measurements in the ultraviolet–visible region showed that the HoF3 contents affected the peak pattern. In that case, the Judd–Ofelt theory [30], used along with the results of the absorption spectra and the refractive indices measurements, can be useful to obtain information about the glass structure. Some reports have described dependence of the spectra on the glass matrix species using calculation of Judd– Ofelt intensity parameters (Ol parameters) for Ho3þ in different host lattices [12,23, 31– 38]. The refractive indices were measured at 488, 540 and 641 nm to calculate the Ol parameters. They were used to determine the relation between the refractive index (n(l)) and the wavelength (l) by least squares fitting to the Sellmeier’s dispersion equation [38] n2 ðlÞ¼1 þ

Sl2 l2 l0 2

(18:3)

where S and lo are constants. The respective values of S and lo obtained for 10HoF320BaF2-10AlF3-60GeO2 glass were 300 and 1.14. Using Equation (18.3), the refractive index was recalculated at the specific wavelength. On the other hand, no dependence of the refractive index on the wavelength was observed under the condition in this study for 50HoF3-20BaF2-10AlF3-20GeO2 glass. Therefore, n(l) was assumed to be constant, as 1.57. According to the method described in references [12, 23, 30–35, 39–42], the experimental oscillator strengths, fexp, of the aJ -> bJ’ transition at the transition mean wavelength l are presented in Table 18.5 and Table 18.6. From these values, the intensity parameters in the Judd–Ofelt theory were obtained as O2 ¼ 1.96 1020, O4 ¼ 0.64 1020, and O6 ¼ 0.11 1020 cm2 for 10HoF320BaF2-10AlF3-60GeO2 glasses and O2 ¼ 0.05 1020, O4 ¼ 0.10 102 and O6 ¼ 0.02 1020 cm2 for 50HoF3-20BaF2-10AlF3-20GeO2 glasses. The calculated oscillator strength, fcal, is also summarized in Table 18.6. The respective rms deviations of fexp and

Table 18.5 Measured and calculated oscillator strength for Ho3þ ions in 10HoF3-20BaF210AlF3-60GeO2 glass. Reproduced from [2] by permission of Elsevier 5

5

I8 !

F5 S2 5 F4 5 F3 5 F2 3 K8 5 G5 5 G4 3 K7 5

Wavelength (nm)

645.9 542.6 539.0 485.1 474.6 468.4 417.8 384.4 379.3

Oscillator strength f ( x 106) fexp

fcal

1.78 0.388 2.03 0.534 0.438 0.504 0.947 0.587 0.282

1.81 0.634 1.79 0.752 0.300 0.526 1.67 0.187 0.0855

Df ( x 106) 0.030 0.246 0.240 0.218 0.138 0.022 0.723 0.400 0.197

Rare-earth-Containing Oxide Fluoride Glasses

557

Table 18.6 Measured and calculated oscillator strength for Ho3þ ions in 50HoF3-20BaF210AlF3-20GeO2 glass. Reproduced from [2] by permission of Elsevier 5

I8 !

5

F5 S2 5 F4 5 F3 5 F2 3 K8 5 G5 5 G4 3 K7 3 F2 3 F4 5

Wavelength (nm)

645.9 542.6 539.0 485.1 474.6 468.4 417.8 385.4 382.3 357.5 334.3

Oscillator strength f ( x 106) fexp

fcal

0.300 0.117 0.190 0.125 0.0782 0.0275 0.189 0.0151 0.0366 0.258 0.0635

0.278 0.0494 0.276 0.0602 0.0478 0.0331 0.0869 0.0411 0.00861 0.00136 0.00323

Df ( x 106) 0.022 0.068 0.086 0.065 0.030 0.006 0.102 0.026 0.028 0.257 0.060

fcal were d ¼ 3.9 107 and d ¼ 1.1 107 for 10HoF3-20BaF2-10AlF3-60GeO2 and 50HoF3-20BaF2-10AlF3-20GeO2 glasses. Some empirical correlations of the intensity parameter and the local structure of the lanthanide ions have been stated in the literature [12,23, 30–35]. In general, O2 increases with the asymmetry of the local structure and the degree of covalency of the lanthanide–ligand bonds, whereas O6 decreases with the degree of covalency. For glass matrices prepared in this study, no clear conclusions are discernible from Ol parameters because of their large uncertainty, especially for O2, but the small value of O2 together with very small O6 for our glasses compared to that for LaAlO3 crystals containing Ho3þ [42] were inferred to result from the lower covalency of the holmium–ligand bonds: the environment around Ho3þ in the glasses prepared in this study could be strongly ionic. This tendency is apparently truer for 50HoF3-20BaF2-10AlF320GeO2 glasses than for 10HoF3-20BaF2- 10AlF3-60GeO2 glass. The ratios of O2/O6 were, respectively, 20 and 2.5 for 10HoF3-20BaF2-10AlF3-60GeO2 and 50HoF3-20BaF210AlF3- 20GeO2 glasses. These ratios imply that the local structure in the 50HoF320BaF2-10AlF3-20GeO2 glasses is more symmetric than that in 10HoF3-20BaF210AlF3-60GeO2 glass. This fact is reflected in the change in the intensity ratio in fluorescence spectra with the Ho3þ contents. It can be applied to control the optical properties of this glass system.

18.3.2

Optical Properties of CeF3-BaF2-AlF3-SiO2 Glasses

The glasses in this system were brown when they were produced, even in inert gas (Ar). The brown colour is attributable to the presence of both Ce3þ and Ce4þ ions that have different energy levels in the glass. The energy transition that takes place between Ce3þ and Ce4þ causes the brown colour. This mixed-valence state of cerium ion resulted from hydrolysis and oxidation at high temperature. It was first found as a result of this study that

558

Functionalized Inorganic Fluorides

colourless or light yellow CeF3-BaF2-AlF3-SiO2 glass was obtained instead of brown glass when it was produced in a CO atmosphere. Oxidization of cerium ion is controlled in this case. The amount of Ce3þ in the glass increased and a very low content of Ce4þ was achieved. These glasses show blue emission from Ce3þ under UV irradiation (365 nm). Although previous reports describe that the addition of carbon powder during melting is effective to avoid oxidation of Ce3þ to Ce4þ in the case of oxide glasses, the decomposition and/or hydrolysis of CeF3 was not avoided in the case of the oxide fluoride glasses. In addition, the Pt or Pt/Au container was badly damaged when using a H2 atmosphere. In this work, glasses having different characteristics can be prepared under an Ar or CO atmosphere. The brown Ce3þ - Ce4þ oxide fluoride glass produced in Ar atmosphere emitted no fluorescence, although the glass produced in CO gas exhibited blue emission under UV irradiation (365 nm). Two broad peaks are observed at 320 nm and 348 nm in the excitation spectra for the glass produced in CO gas. The peak at 320 nm was larger than that at 348 nm. Therefore, the wavelength for the excitation of the emission spectra was determined as 320 nm for 10CeF3-20BaF2-10AlF3-60SiO2 glasses. Figure 18.9 depicts profiles of the emission spectra excited at 310, 320, 330, 357, and 366 nm. Radiation absorption and emission of Ce3þ occur because of the transition of an electron from the 4f orbital into the 5d orbital [43]. Generally, the profile in the emission spectrum corresponding to the 4f– 5d transition is broader than in that corresponding to the 4f–4f transition. Figure 18.9 shows that the peak position for excitation at 320 nm was located at a shorter wavelength than that at 357 nm excitation. To investigate these differences in the emission profiles, the peak was deconvoluted using a Gaussian function. Figure 18.10 depicts the result of the peak analysis of the emission profile excited at 320 nm. The results for all profiles portrayed in Figure 18.9 are presented in Table 18.7. Every emission peak consisted of three peaks that have a peak position at about 410 (peak 1), 445 (peak 2), and 490 nm (peak 3). Matsui et al. [44] also reported the existence of three peaks from peak analysis of the Ce3þ emission spectrum in Y2SiO5:Ce3þ crystal. The difference between peak 1 and peak 2 might correspond to the split (2000 cm1) of 2F5/2 and 2F7/2 by spin–orbit interaction [45–48]. Peak 3 corresponds to the presence of Ce3þ with CN¼7, whereas peak 1 and peak 2 correspond to that with CN ¼ 6 reported in the literature. Therefore, Ce3þ of two kinds might be located in different environments in the considered glass. Peak analysis results show that the intensity of peak 3 tends to be stronger when the excitation wavelength is elongated. The quantum efficiency for peak 3 might depend on the environment around the Ce3þ. Figure 18.11 depicts the energy calculated from each wavelength of peak 1 and peak 2. The energy differences between peak 1 and peak 2 are approximately 1600 cm1; they became small as compared to the energy difference of 2F5/2, and 2F7/2. As reported by Matsui et al., the presence of two kinds of Ce3þ cations in Y2SiO5:Ce3þ seems to cause the change in an energy difference between the 2F5/2 and 2F7/2 levels [44]. The results depicted in Figure 18.5 might be consistent with that inference. Variation must be present in the environment around Ce3þ in the glasses here. Figures 18.12–18.15 present XPS spectra of F1s, O1s, Ce3d, and Ce4d in glasses prepared in an Ar and a CO atmospheres. The F1s peak in the glass prepared in a CO atmosphere was more intense than that in the glass prepared in an Ar atmosphere, as presented in Figure 18.6, and the peak position shifted to a lower energy than that of CeF3. Therefore, less fluoride ion exists in the glass prepared in an Ar atmosphere than in the

Intensity / arb. units

Rare-earth-Containing Oxide Fluoride Glasses

559

λEx = 366 nm 357 nm 330 nm 320 nm 310 nm

370

420

470 520 Wavelength / nm

570

Figure 18.9 Emission spectra of 10CeF3-20BaF2-10AlF3-60SiO2 glass prepared in CO for several excitation wavelengths. Reproduced from [1] by permission of Elsevier

Intensity / arb. units

peak 2

350

peak 3 peak 1

400

450 500 Wavelength / nm

550

600

Figure 18.10 Decomposition of the emission spectrum of 10CeF3-20BaF2-10AlF3- 60SiO2 glass prepared in CO under the excitation wavelength of 320 nm. The solid line shows the observed spectrum. Dashed lines show calculated spectra. Reproduced from [1] by permission of Elsevier

glass prepared in a CO atmosphere. This difference is consistent with the result of quantitative analysis of fluoride ion by XFS measurement, as described previously. The oxidative decomposition of Ce3þ containing intermediate (CeOxFy where 2x þ y ¼ 3) such as CeOF generated by hydrolysis of CeF3 was apparently controlled in a CO atmosphere. The fluoride ion attracts the electron cloud more strongly than the oxide ion. Therefore, the electron cloud seems to approach to fluoride ion in CeOxFy (2x þ y ¼ 3) to a greater degree than in CeF3, which causes the smaller binding energy of F1s electron of the oxide fluoride glass than that of CeF3. A single peak existed in XPS spectra of O1s in the glasses prepared

560

Functionalized Inorganic Fluorides 25 000 peak 1: 5d→4f(2F5/2)

Energy level /cm–1

24 500 24 000 23 500 23 000

peak 2: 5d→4f(2F7/2)

22 500 22 000 21 500 300

320

340

360

Excitation wavelength /nm

Figure 18.11 Energy differences of peak 1 and peak 2 for several excitation wavelengths of 10CeF3-20BaF2-10AlF3-60SiO2 glass prepared in CO. Reproduced from [1] by permission of Elsevier

here, although a double peak was observed in that in CeO2. The peak position for the glasses prepared in a CO or an Ar atmosphere was higher than that for the quartz glass. The reason is that the electron density around O2- might be lowered because the electron cloud is withdrawn toward F via cerium ion. The oxide and fluoride ions in the glass have different electronic states, respectively, from those in simple oxide and fluoride. To elucidate the state of the valence of cerium in the glasses, Ce3d and Ce4d spectra are depicted in Figures 18.14 and 18.15. As references for Ce3þ and Ce4þ, we measured CeF3 and CeO2, respectively. The peak profiles for Ce3d and Ce4d electrons in the glasses prepared in a CO and an Ar atmosphere were similar. However, differences were apparent in their peak positions. Figures 18.14 and 18.15 show that the peaks of the glass prepared in a CO atmosphere located at lower binding energy (about 0.3 eV) compared to those in an Ar atmosphere. The peak position and/or profile in XPS spectra depend on the valence state. The mean valence of cerium ion increases when the oxidative decomposition of CeOxFy proceeds. The ratio of Ce3þ/Ce4þ of the glasses prepared in a CO atmosphere was larger than in glasses prepared in an Ar atmosphere. Therefore, the peaks in the XPS spectra of the glasses prepared in a CO atmosphere have to locate at lower binding energy than the glasses prepared in an Ar atmosphere. In fact, this difference was observed in Figures 18.14 and 18.15. In Ce3d spectra, the peak at 919 eV that appeared in the spectrum of CeO2 was not observed in that of the glasses (Figure 8.14). Furthermore, the peak pattern of the glasses in Figure 8.14 differed completely from that of CeF3 [49–51]. Figure 18.9 portrays the Ce4d spectra of both glasses. The peak near 105 eV corresponds to Si2p. The peaks at 109 (not identified yet), 123 (4d5/2), and 126 (4d3/2) eV that appeared in the spectrum of CeO2 were not observed in spectra of the glasses (Figure 8.14). The peak patterns of the glasses in Figure 8.15 differed completely from that of CeF3. The electronic state of cerium ion must be distinctive for the glass obtained in this study. In other words, the presence of both fluoride and oxide ions might impart a unique electronic state to cerium.

Table 18.7 Analysis of emission spectra of 10CeF3-20BaF2-10AlF3-60SiO2 glass using a Gaussian function for several excitation wavelengths. Reproduced from [1] by permission of Elsevier Peak 1

Peak 2

Peak 3

lEx

Amplitude

Center Wavelength nm

FHWM

Amplitude

Center Wavelength nm

FHWM

Amplitude

Center Wavelength nm

FHWM

310 320 330 357 366

27.94 28.92 2539 1678 12.32

410.34–0.78 410.93–0.67 411.61–0.78 413.90–0.27 418.68–0.27

49.18 49.57 48.12 42.50 32.49

61.33 62.66 60.24 60.32 52.77

446.03–3.29 447.12–3.57 445.71–3.98 442.86–3.41 445.26–1.35

77.30 79.97 77.01 76.19 73.34

26.93 25.07 27.65 33.05 43.40

499.00–19.08 500.51–19.01 495.07–23.74 486.55–17.37 473.75–2.75

122.98 130.60 121.49 126.22 138.96

562

Functionalized Inorganic Fluorides

Intensity / arb. units

glass (Ar)

glass (CO)

CeF3

695

690 Binding energy / eV

685

Figure 18.12 F1s spectra of CeF3 and 10CeF3-20BaF2-10AlF3-60SiO2 glasses prepared in CO and Ar atmospheres. Reproduced from [1] by permission of Elsevier

glass (Ar)

Intensity / arb. units

glass (CO)

Quartz glass

CeO2

538

533 Binding energy / eV

528

Figure 18.13 O1s spectra of CeO2, quartz glass and 10CeF3-20BaF2-10AlF3-60SiO2 glasses prepared in CO and Ar atmospheres. Reproduced from [1] by permission of Elsevier

Rare-earth-Containing Oxide Fluoride Glasses

563

919 eV glass (Ar)

Intensity / arb. units

glass (CO)

925

CeO3

CeF3

915

905 895 Binding energy / eV

885

875

Figure 18.14 Ce3d spectra of CeF3, CeO2 and 10CeF3-20BaF2-10AlF3-60SiO2 glasses prepared in Ar and CO atmospheres. Reproduced from [1] by permission of Elsevier

4d 3/2 4d 5/2

Intensity / arb. units

glass (Ar)

glass (CO)

CeO3

CeF3 130

125

120 115 110 Binding energy / eV

105

100

Figure 18.15 Ce4d spectra of CeF3, CeO2 and 10CeF3-20BaF2-10AlF3-60SiO2 glasses prepared in Ar and CO atmospheres. Reproduced from [1] by permission of Elsevier

564

18.3.3

Functionalized Inorganic Fluorides

Optical Properties of the Glasses Co-doped with TbF3 and SmF3

Reportedly, some rare earth co-doped systems exhibit enhancement of fluorescence intensity [52–54]. Oxide fluoride glasses that can contain rare earth fluorides of more than 50 mol % have been reported above [55, 56]. It is interesting to study the emission properties of such glasses in which two rare earth ions are near one another. This section presents a description of some unique optical properties of the glass containing a large amount of TbF3 doped with a small amount of SmF3 were described. Figure 18.16 presents the relation between I600/540 (the ratio of the peak intensity at 540 nm originated from Tb3þ to that at 600 nm originated from Sm3þ) and the contents of SmF3. In fact, I600/540 increased concomitantly with increasing content of SmF3 (< 2 wt %), whereas I600/540 decreased with the contents of SmF3 (> 2 wt %). Some reports have described that concentration quenching of Sm3þ fluorescence occurred around 1–2 mol % [57,58]. The concentration quenching of Sm3þ fluorescence occurs when the content of SmF3 is more than 2 wt % in Figure 18.16. The intensities of fluorescence at 600 nm of Sm3þ in 20TbF3-20BaF2 – 10AlF3 – 50GeO2 / mol % þ 2 wt % SmF3 were measured as 1.2 104 cps, whereas that in 20TbF3–20BaF2–10AlF3–50GeO2 / mol % þ 5 wt % SmF3 was 5.3 103 cps. This fact reflects that Sm3þ is dispersed identically in the 20TbF3 – 20BaF2 – 10AlF3 – 50GeO2 glass without a phase-separated or otherwise clustered situation. Figure 8.17 presents the relation between I540/600 and the contents of TbF3. Results showed that I540/600 was proportional to x2.05 by least-squares fitting, where x was in xTbF3-20BaF2-10AlF3-(70-x)GeO2 (mol %). The distance of Tb3þTb3þ in the glass (D) is inversely proportional to the third power of the Tb3þ concentration, so I540/600 is almost perfectly inversely proportional to the sixth power of D in this case. The theoretical calculation of Forster–Dexter [59, 60] indicated that the resonant energy transfer probability is inversely proportional to the sixth power of the distance between two centres if the two centres belong to a dipolar transition. Therefore, the change in the relation between I540/600 and the contents of TbF3 presented in Figure 18.17 are explainable mainly using 14.0 Intensity ratio I600/540

12.0 10.0 8.0 6.0 4.0 2.0 0 0

0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0 5.5 x / wt%

Figure 18.16 Relation between the intensity ratio (I540/600) and contents of SmF3 in 20TbF3 – 20BaF2 – 10AlF3 – 50GeO2 /mol% þ x / wt% SmF3 glasses. Reproduced from [2] by permission of Elsevier

Rare-earth-Containing Oxide Fluoride Glasses

565

Intensity ratio I540/600

5.0 4.0 3.0 2.0 1.0 0

0

5

10

15

20 25 x / mol%

30

35

40

45

Figure 18.17 Relation between the intensity ratio (I540/600) and contents of TbF3 in xTbF3 – 20BaF2 – 10AlF3 – (70–x)GeO2 /mol% þ 0.05 / wt% SmF3 glasses. Reproduced from [2] by permission of Elsevier

concentration quenching of the Tb3þ fluorescence [61]. The TbF3 content must be more than 20 mol % to obtain a glass that exhibits an intense orange emission. This phenomenon is a unique property for this LnF3 - BaF2 - AlF3 - GeO2 system that can contain TbF3 more than 50 mol %. The fluorescence of 20TbF3 – 20BaF2 – 10AlF3 – 50GeO2 þ 0.05wt %SmF3 glass exhibited temperature dependence. The intensities of the peak at 540 nm originated from Tb3þ and the peak at 600 nm that originated from Sm3þ are mutually equivalent at room temperature. After heating to 573 K or cooling to 77 K, the intensity of the fluorescence that originated from Sm3þ decreased. The emission colour changed from orange to green through yellow. The glass phase was stable at temperatures of 77–573 K because 673 K is much less than the glass transition temperature (876 K) of 20TbF3 – 20BaF2 – 10AlF3 – 50GeO2 þ 0.05 wt %SmF3 glass. The emission colour changed reversibly according to the temperature. Figure 18.18 portrays the relation between the intensity ratio of I600/540 and temperature. The change in the emission colour from orange to green was recognized when I600/540 was less than 0.7: I600/540 decreased gradually by heating to 673 K. That emission colour changed from orange to green through yellow. In the case of cooling, I600/540 changed drastically and the emission colour changed to green at around 77 K. Relating the fluorescence peak to the energy transition J ! J’, where J and J’ correspond to initial and final states respectively, the electrons in the ground state are excited and the population of J’ becomes higher with increasing temperature. The electrons at J therefore barely transfer to J’. The intensity of the emission was lowered. The Tb3þ has a larger energy gap separating J and J’ than Sm3þ does. Therefore, the temperature affected the emission from Tb3þ only slightly; no significant concentration quenching occurred in the case of Tb3þ [62]. Figure 18.18 shows that, when the temperature rose, the emission from Tb3þ came to take precedence over Sm3þ for the system containing Tb3þ and Sm3þ. For cooling, the mechanism is described as follows. Energy used for the emission from Sm3þ was supplied by nonradiative relaxation from Tb3þ. This nonradiative relaxation is called multiphonon relaxation, which relates to the lattice vibration and dipole-dipole interaction [63,64]. This multiphonon relaxation occurred only slightly around 77 K. Therefore, the energy cannot

566

Functionalized Inorganic Fluorides

be supplied from Tb3þ to Sm3þ when the sample temperature was lower than 77 K. Consequently, the emission from Tb3þ was predominant and the emission colour was green around 77 K. It is expected that these phenomena are useful to probe the energy transfer mechanism in the glass matrix.

Intensity ratio I600/540

1.2 1.0 0.8 0.6 0.4 0.2 0

0

100

200

300 400 Temperature / K

500

600

700

Figure 18.18 Temperature dependence of the intensity ratio (I600/540) of 20TbF3 – 20BaF2 – 10AlF3 – 50GeO2 þ 0.05wt%SmF3 glass. Reproduced from [2] by permission of Elsevier

18.3.4

Magnetic Property of TbF3 Containing Oxide Fluoride Glasses

In fact, Tb3þ is a rare-earth ion having a large effective magnetic moment, 9.72 mB, which is third among the rare earth ions. Considering the transparency of Tb3þ in the visible light region, the Tb3þ-containing compounds are very interesting to prepare material for optical and magnetic applications [22,54, 65, 66]. Using a TbF3-BaF2-AlF3-GeO2 system, glasses containing a large amount of Tb3þ (50 mol % as TbF3) were prepared as described above. Magnetic measurements were conducted at 77–273 K using a vibrating sample magnetometer (VSM) (VSM-3; Toei Industry Co., Ltd.) with a maximum magnetic field of 10000 – 10000 Oe. A Physical Properties Measurement System (PPMS) susceptometer (Quantum Design Co.) is also used to investigate the relation between magnetization and temperature at 2–300 K. Figure 18.19 portrays the relation between TbF3 contents in xTbF3-20BaF2-10AlF3-(70-x)GeO2 / mol % glass and magnetic susceptibility at room temperature. Generally, the relation between magnetization and the magnetic field is shown as !

!

M ¼H

(18:4)

where w represents the magnetic susceptibility (emu mol1). The magnetic susceptibility is expressed as Equation (18.5) from the Curie–Weiss laws c ¼ (18:5) T

Rare-earth-Containing Oxide Fluoride Glasses

567

where is the Weiss temperature and C is the Curie constant; C is written as C¼

ng2 m2g J ðJ þ 1Þ nMaf f 2 ¼ 3kB 3kb

(18:6)

where n, g, mB, kB, J, and Meff respectively represent the number of magnetic ions per mol, the Lande´ g-factor, Bohr magneton, total angular momentum, and the effective magnetic moment. The solid line in Figure 18.19 was calculated from Equations (18.4)(18.6). The magnetic susceptibilities are proportional to the contents of TbF3 in the glasses obtained in this study. Their values fit the calculated values. Figure 18.20 portrays the magnetic susceptibility and the reciprocal magnetic susceptibility (w1) of 30TbF3-20BaF2-10AlF3-40GeO2 glass of 2–300 K. The relation between w1 and temperature is linear, as presented in Figure 18.20. The Weiss temperature – which is an intersection point of the w1 axis and temperature axis – is almost zero, meaning that the spin–spin interaction was negligible at 100275 K in the glasses obtained here. From Equations (18.2) and (18.3), the Curie constant (C) and the effective magnetic moment (Meff) were calculated. These values are presented in Table 18.8. The Curie constant increased concomitantly with increasing glass TbF3 contents. In addition, the effective magnetic susceptibility is consistent with the theoretical value of 9.78 mB. It is therefore readily apparent that the magnetic moment is derived only from a trivalent terbium ion. The atomic content of Tb3þ in 40TbF3–20BaF2–10AlF3–30GeO2 was calculated as 11 %, which is comparable to that reported for oxide glasses as the maximum one, 12 % [67]. In addition, the saturation behaviour might result from the Tb3þ(#)–O2–Tb3þ(") superexchange interaction, which prevented orientation of Tb3þ magnetic moments to the applied magnetic fields in the case of the oxide glasses [68]. Using fluoride instead of oxide is one way to restrain this saturation behaviour. The atomic content of Tb3þ in 50TbF3–20BaF2– 10AlF3–20GeO2 glass obtained in this study was calculated as 14 %. Therefore, it might have high potential for use as a material for Faraday devices [69].

Magnetic susceptibility χ / emu mol–1

0.02

0.015

0.01

0.005

0

10

20 30 x / mol%

40

50

Figure 18.19 Relationship between TbF3 contents in the glass and magnetic susceptibility at room temperature. x is in xTbF3-20BaF2-10AlF3-(70-x)GeO2 (mol%). The solid line represents the theoretical relationship. Reproduced from [2] by permission of Elsevier

Functionalized Inorganic Fluorides 1.0

100

0.8

80

0.6

60

0.4

40

0.2

0

50

100 150 Temperature / K

200

χ–1 / mol emu–1

χ / emu mol–1

568

20 250

Figure 18.20 Temperature dependence of magnetic susceptibility of 30TbF3-20BaF2-10AlF340GeO2 glass. Reproduced from [2] by permission of Elsevier

Table 18.8 Curie constant and effective magnetic susceptibility of xTbF3-20BaF2-10AlF3- (70x)GeO2 glasses (x ¼ 10-40 / mol%). Reproduced from [1] by permission of Elsevier Concentration of TbF3/ mol% Curie constant / K Effective magnetic 10.0 20.0 27.5 30.0 40.0

18.4

1.00 2.47 3.12 4.49 5.13

8.83 9.77 9.37 10.7 9.96

Conclusion

Rare-earth-containing oxide fluoride glasses LnF3 (Ln: Y through Lu)-BaF2-AlF3-GeO2 (or SiO2) were produced in which the nominal content of LnF3 reached 60 mol % maximum. Their basic properties, such as density, refractive index, and glass transition temperature were investigated and summarized in detail. A CO atmosphere is effective to prepare glasses containing a trivalent ion, whose valency might change during the preparation process, such as Ce3þ/Ceþ4. In particular, to discuss the local structure surrounding the rare-earth ion in the glass, a Judd–Ofelt analysis (discussion with O parameters) of the HoF3-BaF2-AlF3-GeO2 glasses was conducted. The fluorescent behaviour and the magnetic properties of LnF3-BaF2-AlF3-GeO2 glasses (Ln ¼ Tb and/or Sm) were also studied to characterize the glasses. Their magnetic and optical properties are attractive for some applications. This glass system has much compositional variety. It might be interesting for applications and for fundamental studies of the lanthanides’ optical properties.

Rare-earth-Containing Oxide Fluoride Glasses

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19 Switchable Hydrophobic-hydrophilic Fluorinated Layer for Offset Processing Alain Tressaud*, Christine Labruge`re and Etienne Durand Institute of Condensed Matter Chemistry of Bordeaux (ICMCB-CNRS), University Bordeaux 1, 87 Avenue du Dr A. Schweitzer, 33608 Pessac, France.

19.1

Introduction

The use of heat-sensitive printing plate precursors has become very popular as one of the ‘dry’ lithographic methods that have recently been proposed to comply with environmental requirements, A major problem associated with most ablative plates, however, is the generation of ablation particles that may contaminate the electronics and optics of the device. The objective of the present study is to provide a ‘processless’ method, through a plasma-enhanced fluorination (PEF) treatment involving a fluorinated gas [1], yielding a positive-working and heat-sensitive layer onto the anodized alumina support. The influence of the radio-frequency (rf) plasma experimental parameters has been investigated with several fluorinated gases. However the discussion will be focused on a comparison between CF4 and c-C4F8 (octafluorocyclobutane) treatments, which clearly illustrates the different possible types of surface reaction – reactive etching and deposition. The hydrophobic properties have been deduced from contact angle measurements and the stability of the formed layer with regard to washing or heating has been also investigated.

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

572

19.2

Functionalized Inorganic Fluorides

The Principles of the Lithographic Printing Process

Lithographic printing methods use a printing master, which is a printing plate mounted on a cylinder of a printing press. The principles of the offset printing are illustrated in Figure 19.1. The master carries a lithographic image on its surface and a print is obtained by applying ink to the image and then transferring the ink from the master onto a receiver material, which is generally paper. In conventional ‘wet’ lithographic printing, ink and aqueous solutions are supplied to the lithographic image, which contains oleophilic (hydrophobic) areas, which are ink-accepting/water-repelling, as well as oleophobic (hydrophilic) areas, which are ink-repelling/water-accepting. Printing masters are generally obtained by the imagewise exposure and processing of an imaging material called plate precursor. More recently ‘dry’ processes have been developed. In addition to the well-known photosensitive plates, which are suitable for UV contact exposure through a film mask, heat-sensitive printing plate precursors have become very popular these last ten years. Such thermal materials offer the advantage of daylight stability and are especially used in the computer-to-plate method wherein the plate precursor is directly exposed, i.e. without the use of a film mask. The material is exposed to heat or to infrared light, which triggers a physico-chemical process. This process could be ablation, polymerization, insolubilization by cross linking of a polymer, heat-induced solubilization, or particle coagulation of a thermoplastic polymer latex. The principle of the most used thermal Image area:

Non-image area:

Oléophilic and hydrophobic

Oléophobic and hydrophilic

water ink

fountain solution Printing ink with a certain water content Plate cylinder with printing plate

int roller

Image area with droplets film of fountain solution of the fountain solution covering the non-image areas

Figure 19.1

Principles of offset printing

Fluorinated Layer for Offset Processing Hydrophilic Aluminium

Hydrophobic Porous Alumina

Contact angle of the pristine alumina surface

Contact angle after PEF treatment

Contact angle after IR irradiation

Hydrophilic 0°–30°

Hydrophobic > 80°

Hydrophilic 0°–30°

Figure 19.2

573

Reversible hydrophilic/hydrophobic (switchable) surface layer

plates is to form an image by a heat-induced solubility difference in an alkaline developer between exposed and nonexposed areas of the coating. The coating typically comprises an oleophilic binder, e.g. a phenolic resin, of which the rate of dissolution in the developer is either reduced (negative working) or increased (positive working) by the image-wise exposure. During processing, the solubility differential leads to the removal of the nonimage/nonprinting areas of the coating, thereby revealing the hydrophilic support, whereas image/printing areas of the coating remain on the support. In this latter process, hydrophilic/oleophobic zones are converted into hydrophobic/oleophilic ones, so that a stronger affinity towards ink is created at the modified areas with respect to the surface of the unexposed areas. A reversible process, using heat pulses, allows the conversion of hydrophobic zones back into hydrophilic ones. A major problem associated with most ablative plates, however, is the generation of ablation particles, which may contaminate the electronics and optics of the device and may also interfere during the printing process. These particles need to be removed from the plate by wiping it with a cleaning solvent, so that ablative plates are often not truly ‘processless’. One of the important issues is therefore to obtain real processless methods.

19.3 19.3.1

Experimental Part Fluorination by Cold rf Plasmas

Radio frequency plasma fluorination is a low-temperature process where fluorinated gases are excited by an rf source and dissociated into chemically active atoms, radicals and molecules. Several fluorinated gases were used: CF4, C3F8 and c-C4F8, which were excited by a rf source at 13.56 MHz. A primary vacuum was obtained by a 40 m3.h1 pump equipped with a liquid nitrogen condenser, which trapped the residual gases. The reactor comprised two cylindrical barrel-type aluminium electrodes which were coated with alumina and which were located within a distance of 2 cm from each other and several gas inlets allowing the use of gas mixtures. The inner electrode was connected to the rf

574

Functionalized Inorganic Fluorides

source, the outer electrode was grounded. The sample to be treated was placed at the centre of the chamber, onto the inner electrode. The gas was introduced in the inner part of the reactor and then dissociated by electron impacts occurring between the two electrodes. Neutral species and radicals diffused from this plasma zone to the centre of the reactor where they reacted with the sample. The sample is generally pretreated with an O2 plasma before the plasma fluorination process, in order to render its surface more reactive. This pretreatment removed adsorbed airborne organic pollution and filled oxygen vacancies at the surface of the sample. Parameters that may be varied during the plasma fluorination process are the pressure and flow of the gases, the temperature and the reaction time. The pressure of the gas in the reactor could be varied between 10 and 300 mTorr, the temperature was thermostatically controlled and maintained at room temperature or a temperature up to 90 C and the duration of the treatment could take from a few minutes up to one hour. All plasma processes were carried out with an rf power of 80 W. The stability of the formed coating layer versus water was tested by washing with distilled water. The treated substrates were composed of a porous, 5 mm-thick, alumina layer elaborated onto an alumina plate of 0.3 mm. These plates, provided by AGFA-Gevaert, were protected with a gum layer. Before using the plate, this layer was removed with rinseage with distilled water and dried by compressed air. In order to characterize the treated plates, two main techniques were used: contact angles measurements and X-ray photoelectron spectroscopy (XPS or ESCA).

19.3.2

Wettability Measurements

Distilled water was used to determine the hydrophilic or hydrophobic aspect of treated alumina plates. The contact angle is described as the angle between the tangent of the droplet at the contact point with the solid and the surface of this droplet. The equilibrium between the liquid droplet and a solid surface in contact with the liquid vapour is expressed by the following Young relation: SV ¼ SL þ L cos, where SV is the surface energy of the solid in contact with the liquid vapour, SL is the interfacial energy between the solid and the liquid, L is the superficial tension of the liquid, is the contact angle of the liquid on the solid. The contact angle is related with the surface tension of the liquid and the surface free energy of the substrate. Complete wettability is achieved when the surface free energy of the substrate is higher or equal to the surface tension of the liquid. In this case, the contact angle is near zero, the surface being termed hydrophilic. On the other hand, the contact angle is very high when the substrate exhibits a low surface energy: the surface is then termed hydrophobic.

19.3.3

Surface Analyses

XPS analyses were performed with a VG 220 i-XL ESCALAB. The radiation was an Mg nonmonochromatized source (1253.6 eV) at 200 W; 200–300 mm diameter areas were investigated on each sample. Surveys and high-resolution spectra were recorded, then fitted with an Eclipse processing program provided by Vacuum Generators. Each C1s component was considered as having similar full width at half maximum (FWHM),

Fluorinated Layer for Offset Processing

575

i.e. 1.2–1.3eV. This choice appears to be in good agreement with our experimental conditions. For instance, a C 1s spectrum taken with a pass energy of 20 eV yields for a monochromatized Al K source components with 1.0 eV FWHM – and even less – and with 1.2 to 1.5 eV for a nonmonochromatized Mg K source. FWHM of F1s components have been fixed at 1.70 eV. We have used a fitting procedure that had been applied with success to several types of carbon-based materials, containing various types of C-F bonds [2,3]. A good agreement between the experimental curve and the full calculated envelope was generally obtained. Despite an important number of required components, this procedure allowed two major phenomena to be taken into account: a primary effect of electron withdrawing of F directly linked to C and a secondary one due to the inductive effect on C nondirectly bound to F with F atoms as nearest neighbours [4].

19.4 19.4.1

Various Types of Surface Modifications using Fluorinated rf Plasmas Reactive Etching of Porous Alumina using CF4-Plasma Treatment

Prior to plasma treatments, the porous alumina substrate exhibits hydrophilic properties, the contact angles of the substrate ranging between 18 and 22. The analysis of the alumina layer reveals the presence of some mineral impurities (Si, Ca, Na) and some contamination elements such as carbon. These impurities could arise from the electrolytic bath used to anodize the aluminium plates and from remainders of the gum coating. During the XPS analysis the plate was etched to 25 nm deep in order to remove the carbon contamination. The Al2p aluminium peak, located at the binding energy (BE): BE ¼ 74.6 eV, corresponds to Al-O bonding as generally found in alumina. In the following such a BE will be taken as an Al-O reference. The surface composition has been deduced from XPS data using Scofield corrections. They are grouped in Table 19.1 for a material treated in the following conditions: O2 pretreatment; p CF4 ¼ 200 mTorr, T ¼ 25 C, t ¼ 60 min. The fitted high resolution Al2p XPS spectrum is given in Figure 19.3. The Al2p envelope can be fitted into three components: the one at lowest BE: 74.2 eV can be assigned to nonfluorinated alumina. The two other components can be attributed to two types of Al-F bondings. This assumption is corroborated by the two components of the fitted high resolution F1s peak in Figure 19.4a. Such an assignment can be proposed with respect to a previous paper by Rodriguez et al. dealing with an XPS study of alumina surface modified by impregnation of NH4F solution [5]. In this study, it was argued that Table 19.1 Surface composition of porous alumina treated under various PEF conditions. Ratios of C-C and C-Fn bonds (Eb 287 eV) Fluorinated medium C (at.%) F (at.%) Al (at.%) O (at.%) Other (at.%) C-Fn(%) C-C(%) CF4 C3F8 c-C4F8

14.8 19.9 32.3

15.8 35.9 66.6

26.4 17.1 –

42.5 25.9 1.0

Ca:0.6 N:1.2 –

0 52 73

100 48 27

576

Functionalized Inorganic Fluorides Al-F, O

Counts / s

Al-O Al-F

82

81

80

79

78 77 76 75 74 Binding Energy (eV)

73

72

71

70

Figure 19.3 High resolution Al2p XPS spectrum of CF4 PEF-treated sample. Reproduced by permission of Science China, Ser. E-Technolog. Sciences 52, pp. 104–110, Springer (2009) Copyright Science in China Press Table 19.2 Element Atomic %

Elemental surface analysis of an a-AlF3 single crystal from XPS analysis F 1s 64.2

O 1s 4.4

C 1s 10.1

Al 2p 21.3

the fitting of the Al2p spectra corresponded to the sum of a peak at 74.3 eV (Al belonging to Al2O3) and of a peak at about 77 eV, which was attributed to Al in a fluorinated environment. The latter component has been assigned to F-Al-F bonds by comparison with two types of pure AlF3 samples: a powdered sample and a single crystal. The structure of both samples was ascribed to the rhombohedral -form [6]. In both cases, the Al2p peak is symmetric and located at 77 eV, whereas that of F1s is located at 686.9 eV. The elemental analysis of the surface of -AlF3 single crystal, given in Table 19.2, leads to a F/Al ratio in good agreement with the composition (F/Alexp ¼ 3.01). The assignment of the components of the F1s envelope is rather complex, because of the occurrence of different sources of F-M bonding. Besides the F-Al-F ones, F-M components may arise from the presence of inorganic impurities such as Ca, Na, Si. On the other hand, as far as the substrate is formed of alumina, F-Al-O bonding are also expected. The component at lowest BE (684.3 eV) is similar to those of F-Ca and F-Na in inorganic fluorides systems. On the other side, the component at 687 eV can be thoroughly assigned to F-Al-F bonding, as found in AlF3 for instance. The main F1s component located at 685.7 eV can be assigned to the presence of F-Al-O bonds. Concerning the C1s spectrum (Figure 19.4), the main peak is located at 284.6 eV and corresponds to C-C bonds of contamination carbon. This peak will serve as reference for measuring the range of the chemical shifts. On the high energy side a shoulder is found at 286.3 eV. This contribution may be due mostly to carboxyl groups. It can be asserted that, using CF4 in PEF conditions, no C-F bonding, generally located at BE higher than 287 eV, can be formed. Table 19.3 sums up the assignment of the different components of Al2p and F1s peaks of CF4 PEF-treated alumina.

Fluorinated Layer for Offset Processing 290.9

577

687.5 686.4

(c)

291.8 289.1288.1

287.1 286.2 285.2

293.8

296

294

Counts / s

Counts / s

289.8 292.8

292

290

288

286

284

282

280

688.8 685.1

692 691 690 689 688 687 686 685 684 683 682 681 680 Binding Energy (eV)

Binding Energy (eV) Csp3

(b)

687.2

285.7

685.8

296

Counts / s

Counts / s

283.6 286.5 287.5 291.3 290.3 289.1 292.4 293.7

294

292

290

288

286

284

282

280

688.6

684.2

692 691 690 689 688 687 686 685 684 683 682 681 680

Binding Energy (eV)

Binding Energy (eV) Csp3

(a)

685.7 684.3

296

Counts / s

Counts / s

286.7 283.7

286.4

287.4

294

292

290

288

286

284

282

280

687.0

692 691 690 689 688 687 686 685 684 683 682 681 680

Binding Energy (eV)

Binding Energy (eV)

Figure 19.4 Fitted high-resolution C1s and F1s XPS spectra of CF4a), C3F8b) and c-C4F8c) PEF-treated porous alumina. Reproduced by permission of Science China, Ser. E-Technolog. Sciences 52, pp. 104–110, Springer (2009) Copyright Science in China Press Table 19.3 Assignment of the different components of Al2p and F1s peaks of CF4 PEF-treated sample, together with Al2O3 and AlF3 references Material

Type of bond

CF4 PEF-treated alumina

O-Al-O F-Al-O F-Al-F O-Al-O F-Al-F

Al2O3 -AlF3

F1s components (eV)

Al2p components (eV)

– 685.7 687.0 – 686.9

74.2 75.3 76.8 74.6 77.3

Contact angles measured on plates just after CF4 plasma treatment were, in most cases, lower those of reference plates before plasma. After washing with distilled water and drying with compressed air, these contact angles improved, and ranged between 130 and 145. The contact angles are given in Table 19.4 with their standard deviations, according

578

Functionalized Inorganic Fluorides

Table 19.4 PEF experimental parameters and contact angles of water on alumina plates after CF4 according to the different post-plasma treatments PEF parameters

Contact angles () After plasma þ washing

p(mTorr) 200 200 200 200

T (C)

Duration (min)

25 90 25 90

60 60 15 15

After plasma þ washing þ drying at 100C

Average value Average value (with standard deviation) (with standard deviation) 101(10) 141(3) 103(8) 134(2) 60(2) 133(2) 87(4) 143(7)

to the different postplasma treatments. No general conclusions can be drawn from these results. Although the contact angles are improved by washing, the results are often inhomogeneous. The more salient effect of washing is a decrease of about 70% of the surface fluorine atomic percentage. This trend could be explained by the reaction that takes place during the fluorination of the sample by the CF4 plasma. A reactive etching occurs on the alumina surface, which is due to the presence in the plasma of a large number of highly reactive F• radicals resulting from the rf decomposition of CF4. This reaction yields the formation on the outermost surface of the plate of very reactive AlF3 particles, which immediately react with ambient moisture to give rise to stable trihydrated aluminium fluoride, AlF3.3H2O. So, the very low contact angles measured just after the plasma process could be explained by the formation onto the surface of a hydrophilic AlF3-based layer. During the drying process, this compound dehydrates and crystallizes into the stable rhombohedral form. Various experimental conditions have been checked in order to approach the involved reaction process. During the XPS analysis, some samples have been etched deep to 300 nm using an Ar beam and have revealed that fluorine is present within the analysed thickness. The concentration decreased for the first 100 nm, then is equal to about 15 %. It seems hazardous to correlate the contact angle variations with any particular plasma parameters. However, considering the fluorine atomic percentage from XPS analysis, it can be argued that surface fluorine increases with the decrease of the substrate temperature (from 90 to 25 C), and with the increase of total pressure (from 50 to 200 mTorr). 19.4.2

Switchable Hydrophilic/Hydrophobic Fluorocarbon Layer Obtained on Porous Alumina using c-C4F8 Plasma Treatment

The C1s and F1s spectra of porous alumina treated with a c-C4F8 rf-plasma are given in Figure 19.4c. The corresponding experimental conditions are: O2 pretreatment; p C4F8 ¼ 200 mTorr, T ¼ 90 C, t ¼ 60 min. The C1s envelope is spread over 10 eV. In addition to contamination sp3C similar to the one found in CF4-treated materials, numerous contributions appear at higher BE, which correspond to C-Fn bonds. The position of

Fluorinated Layer for Offset Processing Table 19.5

579

Assignment of the XPS C1s components of fluorinated carbon

Chemical bond involved sp3C Non-fluorinated C in a fluorinated environment CF in a weakly fluorinated environment CF in fluorinated environments CF2 in a weakly fluorinated environment CF2 in fluorinated environments CF3 and highly fluorinated environments

Binding energy (eV) 284–285 285.5–286.5 287–288 288–289 290–291 291–292.5 293–294

these various C1s components have been determined with respect to saturated non-functionalized sp3 C whose BE after charging-effect correction is set in the 285 eV range. The surface analysis and the assignment of these C-Fn components are given in Table 19.1 and 19.5, respectively. They are in agreement with those previously proposed [2,3]. Above BE ¼ 287.0 eV, the components can be assigned to C atoms that are directly bound to F atoms. The BE increases with increasing number of neighbouring fluorine atoms. In between 285 and 287 eV, the components correspond to C atoms that are not directly bound to F atoms. In this range, the BE shift is due to an inductive effect, which is dependent on the number of F atoms in the position of a given C atom, that are bound to his first C neighbour. The shift can be evaluated to about 0.6 – 0.2 eV for each F atom and is approximately additive. The main C1s peak is located at 290.9 eV and corresponds to C atoms bound to two F atoms. The general shape of the F1s envelope is more symmetrical than those of CF4 and C3F8 treatments and can be fitted into four components. The envelope is centred on F-C binding energy at 687.5 eV. In the F1s spectrum, it is not easy to precisely assign the corresponding peaks, but we can consider that the component at 688.3 eV is due to highly fluorinated C atoms up to CF3, the two components at 687.5 and 686.4 to CF2 groups in more or less fluorinated environments, and the shoulder at lower BE: 685.1 eV to CF groups. In the survey spectra of c-C4F8-treated porous alumina, Al and other atoms (Ca, Na) present in the pristine substrate cannot be detected, which means that the coverage of the surface by the carbon fluoride layer has been achieved. This is not the case for CF4 and C3F8 treatments, as can be noted in Table 19.1. A decrease of the C1s component at higher energies (BE 290 eV), assigned to various C-F bonds, is observed in the depth of the layer, which means that below the surface the fluorocarbon layers disappear. However, in the F1s XPS spectrum, the components do not noticeably decrease and it might be thought that the fluorine atoms tend to be bonded with others elements than carbon in the depth of the layer. Despite the presence of calcium within the layer, which can also be bonded with fluorine to form ionic fluorine at low binding energies, it is clear that the aluminium is bound with fluorine in the interfacial layer between the perfluorocarbon and alumina. The results show that at the interface perfluorocarbon coexists with fluorinated alumina, which explains the occurrence of Al-F-C bonds in this area with binding energy around 685–686 eV. A covalent C-F bond reinforces the ionic character of the Al-F bond and decreases its binding energy. As far as the Ar beam etching does not perturb too much the deposited layer, it can be assumed that two major phenomena take place during the process, as evidenced by the different species

580

Functionalized Inorganic Fluorides

detected at the interface, i.e. a fluorocarbon polymers deposition, accompanied by reactive etching of alumina with formation of Al-F bonds in the subsuperficial zone. The contact angle measured with water shifts from very low values (0 to 30) for the starting hydrophilic surface to above 100 for the c-C4F8 PEF-treated hydrophobic surface. It is important to note that in the case of a c-C4F8 treatment such large contact angles are observed as soon as the treatment is carried out and there is no need to dry the layer as in the case of CF4 plasma. As illustrated by the C1s and F1s XPS spectra, the surface layer is composed of CFn species. It is well known that compounds containing carbon-fluorine bonds such as graphite fluorides, PTFE and other fluorinated polymers are highly stable versus moisture or most chemical reagents. Another specific property is the high level of hydrophobic or water repellency that they exhibit. After this PEF treatment step, the printing plates obtained can be directly mounted on a printing press and a print job can be started without carrying out any processing or rinsing step [7]. During the printing a compressible rubber blanket is used and the prints are made on 80 g offset paper. The ink density was measured on paper by using a GretagMacbeth densitometer Type D19C. The results show that the printing plate based on the untreated aluminium support does not retain ink, with an ink density value £ 0.05, whereas excellent ink-uptake and/or oleophilic properties are observed for a printing plate based on plasmatreated aluminium support, with ink density values 1. After the coating step, the printing plate precursors are dried at 40 C during 30 minutes. Subsequently, the printing plate precursors are irradiated with an IR laser diode, a Nd:YAG or a Nd:YLF laser, at 830 nm with a pitch of 7 qm at varying energy densities, in order to remove the fluorocarbon layer deposited by c-C4F8 plasma treatment. After exposure to infrared light, the printing plates have excellent hydrophilic properties and do not retain ink, as indicated by the low ink density values, with ink density values £ 0.05.

19.5

Comparison of Surface Modifications of Porous Alumina using Various Fluorinated Media: CF4, C3F8 and c-C4F8

The comparison between PEF treatments using various fluorinated gases clearly illustrates the versatility of such a technique. When CF4 is used, the formation during the plasma reaction of a large number of reactive F• radicals give rise mostly to a reactive etching, with the formation at the surface of inorganic fluorides (mostly AlF3). Such a process has been also investigated with other fluorinated gases [8].These species being very sensitive to moisture, a drying of the plate is needed to ensure a higher contact angle. On the other side, octafluorocyclobutane (c-C4F8) is known to give rise to deposition processes because the dominant ion species to be formed during the rf plasma process is C2F4þ. An intermediate behaviour is expected for the octafluoropropane molecule (C3F8) because of its linear conformation,. The comparison between C1s and F1s spectra clearly illustrates the evolution between these three types of PEF treatments (Figure 19.4). The ratio between the nonfluorinated carbon (for BE < 287eV) and fluorinated CFn groups (for BE 287eV) given in Table 19.1 clearly illustrates the evolution in the surface composition ranging from inorganic fluorides to perfluorocarbon. The presence of

Fluorinated Layer for Offset Processing

581

a fluorocarbon surface layer, which is absent in the case of CF4-treated alumina, increases from C3F8 to c-C4F8 treatments. In the case of C3F8 fluorinated sample, an inhomogeneity of the layer composition is noticeable on both F1s and C1s spectra. New components at a high BE are indeed present (at 688.6 and 293.7 eV), which account for the presence of overfluorinated -CF3 groups with respect to CF2 and CF groups, which constitute most species of the c-C4F8 fluorinated sample. Moreover, for similar deposition conditions, it is necessary to etch the c-C4F8 fluorinated layer upon 10 nm before observing the presence of aluminium, whereas this element is observed during the surface analysis of the C3F8 fluorinated sample. It is clear that the polymerized layer is thicker using the c-C4F8 than the C3F8 medium and leads to more homogeneous and less crosslinked layers.

19.6

Conclusion

In conclusion, by rf-plasma fluorination the outmost surface of the hydrophilic anodized aluminium support is hydrophobized, or in other words, the hydrophilic support is converted into a support with hydrophobic properties, as shown in Figure 19.5. This conversion from a hydrophilic to a hydrophobic state is characterized by an increase of the contact angle measured on the surface, indicating an increase in the hydrophobic-to-hydrophilic ratio. The most adapted route to drastically switch from a pristine hydrophilic surface of porous alumina into a hydrophobic layer is clearly the c-C4F8 treatment. It has been shown

γ S V = γ S L + γ L cos θ γL γSV

θ

γSL

γ SV: surface energy solid/liquid vapour γ SL: interfacial energy solid/liquid γ L: superficial tension of the liquid θ: contact angle of the liquid on the solid

substrate

Contact angle θ

Untreated substrate

CF4 rf plasma

≈ 20°

≈ 60°

hydrophilic

C4 F8 rf plasma

≥ 130°

hydrophobic

Figure 19.5 Contact angles of pristine and plasma-treated porous alumina. Reproduced by permission of Science China, Ser. E-Technolog. Sciences 52, pp. 104–110, Springer (2009) Copyright Science in China Press

582

Functionalized Inorganic Fluorides

that the formed hydrophobic layer can be further switched back to a hydrophilic one through its vaporization using a laser beam. The fluorinated layer can be evaporated during the heating process so no ablation particles are retained on the material.

Acknowledgements These investigations have been carried out with the support of AGFA GEVAERT NV, Septestraat, 27, 2640 Mortsel, Belgium. Dr J. P. Chaminade is acknowledged for the growth of AlF3 single crystals and Dr A. Demourgues for fruitful discussions and suggestions. Contact angles measurements were carried out at Biomate´riaux et re´paration tissulaire, INSERM-U577, Universite´ Victor Segalen Bordeaux 2, France.

References [1] C. Cardinaud, A. Tressaud, ‘Surface modification of inorganic materials by fluorination’, pp. 437–492, in Advanced Inorganic Fluorides, T. Nakajima, A. Tressaud, B. Zˇemva, Eds. Elsevier, Amsterdam (2000). [2] G. Nanse´, E. Papirer, P. Fioux, F. Moguet, A. Tressaud, ‘Fluorination of carbon blacks’, Carbon 35, pp. 175–194, 371–388 and 515–528 (1997). [3] A. Tressaud, E. Durand, C. Labruge`re, ‘Plasma-Enhanced Fluorination of Nitrile Butadiene Elastomer: an XPS study’, in Plasma Processes and Polymers, R. d’Agostino, P. Favia Eds. Wiley-VCH Verlag GmbH, Weinheim, (2005). [4] A. Tressaud, E. Durand, C. Labruge`re, ‘Surface modification of carbon-based materials: comparison between CF4 rf cold plasma and direct F2-gas fluorination routes’, J. Fluorine Chem., 125, 1639–1648 (2004). [5] Rodriguez L. M., Alcaraz J., Hernandez M., Taarit Y. B., Vrinat M. ‘Alkylation of benzene with propylene catalyzed by fluorinated alumina’. Appl. Catal. A: General, 169, pp. 15–27 (1998). [6] D. Babel, A. Tressaud, Chap. 3, ‘Crystal chemistry of fluorides’ in Inorganic Solid Fluorides, P. Hagenmuller Ed., Academic Press, New York (1985). [7] H. Andriessen, C. Brigouleix, A. Tressaud, ‘Processless lithographic printing plate’, European Patent EP 1 640 175 A1, Agfa-Gevaert, March 2006. [8] T. Kawabe, M. Fuyama, S. Narishige, ‘Selective ion beam etching of Al2O3 films’, J. Electochem. Soc. 138, 2744–2748 (1991).

Index

Note: page numbers in italics refer to figures or tables. Acenaphthylene 482–3 Acidity, see Surface acidity studies Acylation reactions 75–6 Adsorption, see Molecular adsorption; Probe molecules Alkoxysilanes 1–2 Alkylation reactions 75–6 Aluminium industries 205–11 Aluminium brazing, see Fluoroaluminate fluxes Aluminium bromofluoride 158 Aluminium chlorofluoride 107, 109–10, 123, 158 Aluminium fluorides 2 crystallographic forms 42–3, 52–3, 369–70 NMR data 141–75 hydrate 44, 57–9, 58, 159–60 IR spectra 23 surface studies 51, 52 modelling/DFT calculations 180–5 synthesis via thermal decomposition 42 see also HS-AlF3 Aluminium fluoride hydrate 44, 57–9, 58 NMR study 159–60 Aluminium hydroxyfluorides HTB structure 43, 44, 47, 49, 50, 51–7 surface acidity study 117–21, 129–36 NMR study 146–7, 159–60 pyrochlore structure 44, 45–7, 48, 50, 51–7 ReO3-type structure 44, 45, 48 synthesis, see Microwave-assisted synthesis

Aluminium isopropoxide 4, 5–6 NMR characterization 9–11 Aluminium silicon alloys 222–3 Amine cations 348 Ammonia, adsorption of 27, 28, 51, 70, 71 binding energies 185–8 Ammonium fluoride 385, 520 Ammonium hydrogenfluoride 210 Anhydrous hydrogen chloride 107, 112–17, 136 Anti-reflection coatings 34, 307, 314, 317–19 Antimony pentafluoride 109 Azobenzenes 483–4 Barium fluoride 213 HS-BaF2 30 in oxide fluoride glasses 549–52 Benzoylation reactions 75 Binary metal fluorides 29–30 Bismuth oxyfluoride 432–3, 454 Bismuth-based cuprates 406–8 Brazing, see Fluoroaluminate fluxes Brønsted acidity 71 and bi-acidic catalysts 78, 82 versus Lewis acidity 74, 78, 82 molecular acids 101, 108–9 and oxidation catalysts 84, 86 C3F-azo+ ion 472, 476 CAB process, see Fluoroaluminate fluxes Cadmium fluoride 95, 96 and glass-ceramics 281–2

Functionalized Inorganic Fluorides: Synthesis, Characterization & Properties of Nanostructured Solids Ó 2010 John Wiley & Sons, Ltd

Edited by Alain Tressaud

584

Index

Calcium fluoride 32, 33, 35 films 319 fluorite structure 435 and glass-ceramics 281 HS-CaF2 30, 95, 96 single crystal and optical systems 313 synthetic 213 Carbon monoxide, adsorption of 27–8, 70–1, 82, 83, 131–5 hydroxylated surfaces characterized 193–6 Catalysts, HS-metal fluoride 42–3, 69–71, 108 host–guest metal fluoride systems 74–8 HS-AlF3 33–4, 69–74 -AlF3 active sites 196–7 dismutation of CCl2F2 196–201 hydroxy(oxo)fluorides as bi-acidic catalysts 78–83 metal fluoride supported noble metals 32, 88–90 hydrodehalogenation reactions 90–5 Suzuki coupling reactions 95–7 oxidation catalysts 84–8 test reactions 28–9, 112 vanadium oxide 32, 84–8 Cation exchange capacity 472 Cerium-based oxyfluorides, see UV absorbing Ce and Ti-based oxyfluorides CFCs 29, 90, 206 Chemical vapour deposition (CVD) 336 Chlorine-36 label, see Radiotracer studies Citronellal 83 Clays 471, 506–7 Cluster model 435–7 Coatings, see Thin films Colossal magnetoresistance, see Magnetoresistive materials Commercial products, see Industrial materials and applications Complex metal fluorides 30, 43 Complex transition metal oxides 383–4 Composite deposition technology 224 Controlled atmosphere brazing (CAB), see Fluoroaluminate fluxes Copper-based oxyfluorides, see Superconducting cuprates Correlation NMR techniques 148–52, 163 Cross polarization techniques (CP) 148, 150 Cryolite 211 structure 350–1 CRYSTAL code 176 Cuprate superconductors, see Superconducting cuprates

Defect clusters 435–9 Defect region model 443–6 Dehydrochlorination 112, 114 Dehydrofluorination 72–3 Dehydrogenation 84–5, 86 Devitrification 275–7, 287, 288 and optical properties 293–5 DFT calculations, see Modelling/DFT calculations Dichloroacetic acid 94–5 Dichlorodifluoromethane 196–201 Dielectric heating 40–1 Differential thermal analysis 282–7 Dipolar polarization 40–1 Dismutation reactions 29, 72, 75–8, 93 mechanism for CCl2F2 196–201 Doping ions 279, 299–301, 313–17 and ionic conductivity 431–4 Double resonance techniques 148–50 Electrolytes, fluoride 423–7 Electroplating 212 Elpasolite 351 Erbium doping 314, 316 nucleation 284–6, 288–9 Etching processes 206, 210 rf plasma fluorination 571, 573–4 Europium doping 314, 315, 316–17 and luminescence 322–6 Exfoliation 475 Fluorescence dyes 479–82 glass-ceramics 295–8, 343, 344 oxide fluoride glasses 558, 564–5 pyrene 478–9 Fluoride glasses 307, 308–9, 331–3 nanocomposite/thin film materials reviewed 313–17 optical properties interactions with light 313, 314 luminescence 322–6 sol-gel processes and fluorinating agents 309–13 and waveguides, see Waveguides, planar Fluoride-ion mobility, see Ionic conductivity Fluorinating reagents 309–13, 520, 573 Fluorine, elemental 207–8, 214 direct fluorination of silica 520–1 Fluorite, see Calcium fluoride

Fluorite structure 279, 280, 423–4, 426–7, 436 defect clusters/cluster models 435–9 Fluoroalkoxide precursors 311–12 Fluoroaluminate fluxes 210–11, 214 CAB technology 220–1, 224–5 steps in brazing process 220–1 cleaning and flux application 221–2 filler metal alloys 222–3 clad-less brazing 223–4 flux composition 214–16 flux particle size 217–18 hydrogen fluoride generation 216–17 Fluoroaluminates 30, 43, 214 crystallization of, see Polyanion condensation, fluoroaluminate industrial applications 211 see also Fluoroaluminate fluxes for modelling NMR parameters 160–1 Fluoroboric acid 212 Fluorohectorites 508, 510 Fluorolytic sol-gel synthesis coupled with microwave irradiation 47–8 preparation of HS-AlF3 and HS-MgF2 the fluorolytic sol-gel route 4–5 mechanism of fluorination 7–8, 22 optimization of procedure 5–8 single crystal intermediates 13–15 stages in sol-gel fluorination 15–18 structural changes on fluorination 11–13 structures of wet/dry gels compared 18–19 network formation 19–21 properties and structure of HS-AlF3 7–8, 23–9, 69–71, 157–8 range of fluorides obtainable 29–30 see also under Catalysts, HS-metal fluoride; Surface activity suggested applications bifunctional oxide fluorides 31–2 catalysis 32, 33 coatings 33, 34–5 nano-sized metal fluorides 31 optics 34 organically modified/hybrid materials 32–3 Fluorosilicic acid 312 Fluorozirconate glass, see ZBLA(N) glasses Flux, see Fluoroaluminate fluxes Friedel-Crafts reactions 75–6, 108 FTIR spectroscopy 103–7, 110–17, 527–9 Gas insulated switchgear/lines 209 Germanates glass-ceramics 281 microporous materials 499–504

Index 585 Glass, see Fluoride glasses; Oxide fluoride glasses Glass-ceramic systems, oxyfluoride 273–4, 301–2, 323–4 CaF2 as crystalline phase 281 CdF2/PbF2 as crystalline phase 281–2 LaF3 as crystalline phase 282 -PbF2 as crystalline phase 281, 316 differential thermal analysis 282–7 Thakur’s method 288–9 multi-doped systems 299–301 optical properties 293–5, 316–17, 322–4 glass-ceramics compared to single crystals 297 absorption spectra 297–8 spectroscopic properties of Ln3+ excited state’s lifetime 294–6 fluorescence decay 295–6, 298 linewidth effects 294, 294–5 upconversion fluorescence 296–7 preparation and synthesis glass-devitrification processes 275, 315 nucleation/growth 276–7 spinodal decomposition 275–6 sol-gel synthesis 309, 312–13, 315–17 transparency and particle size 277–8 waveguide developments 342–4 Glassy phases 279 Graphite 474 Guest–host, see Host–guest metal fluoride systems Heavy metal fluoride glasses 307, 309, 310–11, 313 Hexafluoroaluminates 2, 201 structures 350–2, 369–70 Hexagonal tungsten bronze structure 43, 51, 180–2, 231, 269, 366, 367 High surface area fluorides, see Aluminium hydroxyfluorides; HS-AlF3 High-temperature superconductors, see Superconducting cuprates Highly fluorinated silica analysis and characterization techniques 521–3 elemental fluorine route 520 fluorination process 520–1, 522 FTIR analysis silanol groups/grafted fluorine correlation 527–9 pore size distribution 529–30 silicon–fluorine bonds 526–7

586

Index

Highly fluorinated silica (Continued ) water adsorption/hydroxylation 523–6 D20 exchange/silanol content 525–6 mechanisms 540–1 nuclear magnetic resonance studies 533 1 H-29Si CP-MAS spectra 534–6, 539 19 29 F- Si CP-MAS spectra 536–9 thermogravimetric analysis/mass spectroscopy 530–1 hydrophobic characteristics 533 thermal stability 530–3 Host–guest metal fluoride systems 30, 74–8 HS-AlF3 characterization IR spectroscopy 23, 24 solid state NMR 24–7, 157–8 surface adsorption/desorption isotherms 7–8 X-ray/TEM 23–4, 25, 51 properties adsorption/Lewis acidity 27–8 catalytic properties 28–9, 69, 71–4, 196–201 surface acidity 27–9, 69–71, 117–21, 123–7 synthesis, see Fluorolytic sol-gel synthesis HS-BaF2 30 HS-CaF2 30, 95, 96 HS-MgF2 5, 30, 69, 70 bi-acidic hydroxy(oxo)fluorides 78–83 host-guest systems 74, 75–8 surface acidity study 120–3 Suzuki coupling 95, 96 HTB hydroxyfluorides, see Aluminium hydroxyfluorides Hybrid materials, see Organic-inorganic hybrid Hydrodehalogenation 90–5 Hydrofluoric acid 206–7 Hydrofluorination 76, 77 Hydrogen chloride probe 107, 112–17, 134 Hydrogen fluoride, anhydrous 205–6 Hydrothermal synthesis 509–10 Hydroxylation, surface 193–6 Hydroxy(oxo)fluorides 78–83 Industrial materials and applications 205 ammonium hydrogenfluoride (bifluoride) 210 barium fluoride 213 calcium fluoride 213 cryolyte and related hexafluoroaluminates 211 fluorine gas 207–8 fluoroboric acid 212

hydrogen fluoride/hydrofluoric acid 205–6 iodine pentafluoride 208 potassium fluoroaluminate 214 see also Fluoroaluminate fluxes potassium fluoroborate 212 potassium fluorometalates 210–11 potassium hydrogenfluoride (bifluoride) 214 sodium fluoride 213 sodium hydrogenfluoride (bifluoride) 213–14 sulfur hexafluoride 209–10 Intercalation compounds, perfluoroalkyl 469–70 compound preparation/synthesis routes 471–5 layered host materials 471, 474 photofunctional molecules fluorescence of rhodamine B 480–2 introduction into layered materials 476–8 micro-environment study 478–80 photochemical properties 482–4 Intersalation 472 Iodine pentafluoride 208 Ionic conductivity 41, 423–7 cation doping of fluorides heterovalent replacement 432 isovalent replacement 431–2 lone electron pair deformation 432–4 fluorite structure/nonstoichiometric phases defect clusters 435–9 fluorite structure 435 micro-ordering/nanostructuring 447–9 ordered fluorite-like phases 39–40 phase diagrams 441, 442 ionic transfer in fluorite-like phases 441, 443 defect region model 443–6 ionic transfer in tysonite-like phases 454–5 concentration dependence 459–62 migration paths in LaF3 structure 455–7 temperature dependence/anion defect positions 457–9 heterovalent oxygen substitution 428–9 oxyfluoride conductivity 429–31 pyrohydrolysis 427 tysonite structure/nonstoichiometric phases 449–50 hexagonal modification 450 ordered phases 454 trigonal modification 450–3 Ionic exchange processes 336 Ionic liquids 64 Ionic transfer activation energy 426

Index Ionothermal synthesis 504 Isomerization 29, 72, 73–4, 112 Isopolegol 83 ITQ-n zeolites 490, 491 Lanthanum cuprate 398, 400–3, 406 Lanthanum fluoride 82 and glass-ceramics 316 host crystal for luminescence 314 ionic conductivity 455–7 Lanthanum oxyfluorides 323–4 see also Glass-ceramic systems Lanthanum strontium cuprates 408–10 Lanthanum strontium manganese oxide 412–13 Lasers 334–6, 341–2 Layered materials, see Intercalation compounds, perfluoroalkyl Lead fluoride 281, 316 Lead tin fluoride 433–5 Lewis acidity of HS-AlF3 27–9, 43, 69, 72 see also Surface acidity studies Lithium hexafluoroaluminate 211 Lithographic methods, see Offset processing Low refractive index films 320–2 Lutidine probe 131–5 Luminescence 307, 314, 315, 316, 317, 322–6 fluorescence glass-ceramics 295–8, 343, 344 photofunctional molecules 478–84 ZELA waveguides 343, 344 Magic angle spinning, see Nuclear magnetic resonance Magnesium fluoride 2, 32, 34 anti-reflective coatings 314 films 319–22 see also HS-MgF2 Magnetoresistive materials 411–15 Manganese-based oxyfluorides 383–4 fluorination of manganites 411–15 MAS NMR, see Nuclear magnetic resonance Mesoporous silica 520–1 Metal oxide acidity 102, 103–7 Metallurgy 209, 210–11, 214 see also Fluoroaluminate fluxes MFI-type zeolite 486, 490–1 Micas 508 Michael addition reactions 78–80 Microporous materials 489–91 germanium-based 499–504 gallo- and alumino-germanates 499–500 ionothermal synthesis 504 materials/OSDA/IZA codes 501–2

587

orienting role of Ge 502–3 pure germanates 500, 502 role of hexamethonium OSDA 503 phosphate-based 504–6 alumino- and gallophosphates 504 role of fluoride 504–6 silica-based 490–9 absence of connectivity defects 497–8 isomorphous aluminium substitution 498–9 ITQ-n zeolites 490, 491 materials/OSDA/IZA codes 492–3 mechanism studies 496–7 MFI-type 489, 490 synthetic clays 506–7 hydrothermal synthesis 509–10 semi-synthetic 507–8 solid state synthesis 508 Microwave-assisted synthesis advantages and principles 39–41 ovens and procedures 41–2 aluminium-based fluorides/ hydroxyfluorides 42–4 reaction parameters/nature of product 44–9 sol-gel alkoxy-fluoride route 47–9 structural features HTB and pyrochlore hydroxyfluorides 51–7 textural features 49–51 transition metal oxy-hydroxy-fluorides 61–3, 232–3 metal organic frameworks 64 tin titanium oxy-hydroxyfluoride 62–3 use of ionic liquids 64 Mixed metal fluoride systems 36, 74–8 Modelling/DFT calculations 175–6, 201–3 density functional theory 176–7 kinetic Monte Carlo simulations 179–80 molecular adsorption 178–9 surface free energies 177–8 characterization of AlF3 surfaces/NH3 binding energies 185–8 geometric structure of - and -AlF3 bulk phases 180 surfaces and surface energies 180–5 and high resolution NMR studies 160–7 surface catalysis 196–7 surface catalysis/CCl2F2 dismutation 196–7 molecular adsorption 197–8 reaction kinetics 200–1 reaction mechanism and barriers 198–200

588

Index

Modelling/DFT calculations (Continued ) surface composition under reaction conditions 188 -AlF3 (0001) termination 192–3 -AlF3-x (01–12) termination 189–91 Molecular adsorption 197–8 calculations 178–9, 188–93 see also Probe molecules Monochlorodifluoromethane 90–4 Monte Carlo simulations 179–80 Montmorillonite 509–10 MQ-MAS 145–6 Multi-doped glass-ceramics 299–301 Multiphonon relaxations 324, 333 Multiple-quantum MAS 145–6 Nano-particulate materials 69–70, 102–3 particle size and transparency 31, 277, 313 see also Aluminium hydroxyfluorides; HS-AlF3 Nanocomposite thin films 313–17 Nanosheet solutions 475 Neodymium cuprate 389–90, 404–5 Neodymium fluoride 384–5 Neodymium oxyfluoride 432–3 Noble metal catalysts, fluoride supported 32, 33, 88–90 NOCOLOKÒ flux, see Fluoroaluminate fluxes Nuclear magnetic resonance 141–2 application of solid-state NMR 27 Al NMR MQ-MAS 145 SATRAS 145, 146 TOP spectra 146–7 19 F NMR/fast MAS and high magnetic field 142–5 correlation 1D/2D double resonance 148–50 2D MAS CP-HECTOR 150–1 2D SLF 150 19 F DQ-SQ MAS 151–2 CP-MAS 148 characterization aluminium fluorides 153 chlorofluoride and bromofluoride 158 fluorinated aluminas and zeolites/ HS-AlF3 157–8 isotropic chemical shift data 153–7 pentahedral and tetrahedral species 158–9 cerium–calcium-based oxyfluorides 247–52 highly fluorinated silica 533–9 HS-aluminium fluoride hydrate 159–60

HS-aluminium hydroxyfluorides 159–60 intermediates in sol-gel synthesis 9–23 modelling/DFT calculations model compounds/fluoroaluminates 160–1 19 F NMR line assignments 161–3 27 Al site assignments 164–7 Nucleation 276–7, 284–6, 288–9, 290

Octafluorocylobutane 571 Offset processing background to lithographic printing 571–3 fluorination and analysis methods rf plasma fluorination of plates 573–4 wettability measurements 574 XPS analysis 574–5 surface modifications/plasma treatments c-C4F8 plasma 578–81 CF4 plasma 575–8, 580–1 switchable hydrophilic/hydrophobic surfaces 573, 580, 581–2 Optical amplifiers 314, 334–6, 340–1 Optical waveguides, see Waveguides, planar Optics overview 2, 34, 307, 308–9 Organic structure-directing agents (OSDA) 490, 491, 492–3 Organic–inorganic hybrids 32–3, 64, 348–9 see also Intercalation compounds, perfluoroalkyl Oswald ripening 277 Oxidation catalysts 32, 84–8 Oxide fluoride glasses 545–6 composition and preparation HoF3-BaF2-AlF3-GeO2 systems 549, 550, 551–2 NdF3-AlF3-GeO2 system 546–7 NdF3-SiO2-Al2O3 system 546 TbF3-AlF3-GeO2 system 547 TbF3-BaF2-GeO2 system 547, 548, 549 density and refractive index 552–3 glass transition temperature 553–5 magnetic properties of TbF3 containing glasses 566–8 optical properties CeF3-BaF2-AlF3-SiO2 glasses 555–7, 557–63 HoF3-BaF2-AlF3-GeO2 glasses 555–7, 557–63 TbF3 and SmF3 co-doped glasses 564–6 Oxyfluorides Ce and Ti-based, see UV absorbing Ce and Ti-based oxyfluorides ionic conductivity 29–31

Index superconducting, see Superconducting cuprates transparent systems, see Glass-ceramic systems, oxyfluoride Palladium catalysts 32, 33, 89–94 Suzuki coupling 95–6 Perfluoroalkyls 208 intercalation compounds, see Intercalation compounds, perfluoroalkyl Photofunctional molecules 469, 476–8 photochemical properties 482–4 photophysical responses/properties 478–82 Photoisomerization 483 Photolithography 338 Photoluminescence, see Luminescence Phyllosilicates 506–7, 509 Physical vapour deposition (PVD) 337 Planar waveguides, see Waveguides, planar Plastic fuel/gas tanks 207 Platinum catalysts 32, 89–94 Polyanion condensation, fluoroaluminate background 347 amine cation templates 348 synthesis 348–9 concentration of amine 372–4 temperature 374–5 extended finite polyanions (0D) 350 Al2F10 dimers 353–4 Al2F11 dimers 353 Al3F16 trimers 353 Al4F18 tetramers 354–5 Al4F20 tetramers 354 Al5F26 pentamers 355 Al7F30 heptamers 355–6 Al8F35 octamers 356 isolated AlF4 tetrahedra 350 isolated AlF6 octahedra 350–2 mixed polyanions 356–8 1D networks Al2F8 double chains 360 Al2F9 chains 359 Al3F12 triple chains 361 Al6F24 ramified chains 361 Al7F29 chains 360 AlF5 chains 358–9 mixed polyanions and/or chains 361–4 2D networks Al2F7 layers 366–7 Al3F10 layers 368 Al3F14 layers 365 AlF4 layers 365–6

589

3D networks Al2F9 368–9 Al7F33 368, 369 AlF3 369–71 Porous metal fluoride coatings 319 Potassium fluoroaluminate 211, 214 see also Fluoroaluminate fluxes Potassium fluoroborate 212 Potassium fluorosilicate 210–11 Potassium fluorozincate 210–11 Potassium hydrogenfluoride 214 Predictive modelling, see Modelling/DFT calculations Probe molecules 27–8, 78, 131–6 anhydrous hydrogen chloride 107, 112–17 Propane, dehydrogenation of 84–5, 86 Protective coatings 34–5 Pulsed laser deposition (PLD) 337 Pyrene 476, 478–9 Pyridine, adsorption of fluorides 27, 28, 70, 81, 131 metal oxides 103–7 Pyrochlore hydroxyfluorides, see Aluminium hydroxyfluorides Pyrohydrolysis 427 Radiotracer studies 110–17 Rare-earths doping 279 emissions and energy levels 334 see also Oxide fluoride glasses REAPDOR technique 148–50 REDOR technique 148–50 Refractive index 34, 320–2 ReO3-type hydroxyfluorides, see Aluminium hydroxyfluorides Rhodamine B 480–2 Rose bengal dye 479–80, 483 Ruddlesden-Popper series 384, 403–4 Satellite transition spectroscopy (SATRAS) 145, 146 Silica-based materials, see Highly fluorinated silica; Microporous materials Silicon oxyfluoride 312–13 Single crystals 297–8 Sodium fluoride 213 Sodium hexafluoroaluminate 211 Sodium hydrogenfluoride 213–14 Sol-gel synthesis, overviews of 1 fluoride glasses 309–13 metal fluorides 2–4, 47–8, 308–9 metal oxides 1–2, 308 see also Fluorolytic sol-gel synthesis

590

Index

Solid electrolytes, see Electrolytes, fluoride Solid solutions 441, 442, 443 Solvay-Fluor products, see Industrial materials and applications Spin casting 337 Spinodal decomposition 275–6 Sputtering 337 Strontium cuprate 398–400 Strontium manganese oxide 411–12 Sulfur hexafluoride 209–10 Supercluster model 437–9 and microordering 447–9 Superconducting cuprates 383–4 complex oxides 383–4 fluorine insertion/fluorination reagents 384–8 fluorination of nonsuperconducting cuprates 408–11 superconductivity–structure relationships 388–9, 415–16 electron-doped conductors 389–90 hole-doped conductors 390–8 structural rearrangements in fluorinated cuprates 398–411 Surface acidity studies 69–71, 101–2 metal oxides compared to fluorides 102–3 FTIR adsorption studies 103–7 molecular acids vs metal fluorides 101, 108–9 solid metal fluorides aluminium chlorofluoride benchmark 109–10, 123 FTIR methodology 103–7, 110–17 radiotracer (36Cl) approach experimental method 112–17 geometries of HCl adsorption 135–6 HS-AlF3 69–72, 123–7 HS-MgF2 69–71, 127–9 HTB aluminium hydroxyfluorides 129–35 Surface activity 128 Surface free energies 177–8 Surface hydroxylation 188, 193–6 Suzuki coupling reactions 95–7 Switchable hydrophobic/hydrophilic layer, see Offset processing Synthesis routes the fluoride route, see Microporous materials see also Microwave-assisted synthesis; Sol-gel synthesis, overviews of Talc 507–8 Tanabe model 30, 74 TEDOR technique 149–50

Telecom applications 334–6 Temperature programmed desorption 27, 69, 70, 71 and calculated binding energies 187–8 FTIR study of metal oxides 103–7 Tetraethylorthosilicate 312 Tetrafluoromethane 571 Thakur’s method 288–9 Thermal decompositon processes 43 Thermogravimetric analysis 260–3, 530–3 Thin films 33, 314 nanosheet solutions 475 optical systems 34, 313–17 anti-reflection/low refractive index 304, 307, 317–18 porous metal fluoride 319 protective coatings 34–5 transparent conducting 313 Tin oxide 313 Tin titanium oxy-hydroxyfluoride 62–3 Titanium oxy-hydroxyfluorides 61–3 Titanium-based oxyfluorides, see UV absorbing Ce and Ti-based oxyfluorides Tocopherol synthesis 80–2 TOP spectra 146–7 TPD, see Temperature programmed desorption Transition metal catalysts 32, 84–8 Transition metal oxides 383–4 Transition metal oxy-hydroxyfluorides 61–4 Transparency 31, 277–8, 313 TRAPDOR 148–50 Trifluoroacetic acid 309, 310–11 Two-dimensional one pulse spectra 146–7 Tysonite structure 423–4, 426–7, 449–53 Ultra-low dielectric constant 321 Ultra-violet absorption, see UV absorbing Ce and Ti-based oxyfluorides Upconversion processes 296–7, 333 Uranium fluoride 206 UV absorbing Ce and Ti-based oxyfluorides 229–30 analysis and magnetic measurements 233–4 cerium–calcium-based oxyfluorides structure and fluorine environment 237–42 19 F NMR analysis/local fluorine environment 247–52 CaF2 observation/ion solubility limits 242–4 evaluation of cation–anion bond lengths 244–7 TEM and EELS analysis 252, 253

synthesis 231 ultraviolet shielding properties 263–6 titanium-based oxyfluorides structure and composition 252–3 characterization and features 253–60 thermal behaviour 260–3 synthesis 232–3, 252–3 ultraviolet shielding properties 263, 266–7 Vanadium oxide catalysts 32, 84–8 Waveguides, planar 314, 316, 331–2 applications 334–6 fabrication processes 336–8 fluoride glass

Index 591 composition and properties 332, 333 emerging glass-ceramic systems 342–4 geometry and performance 338–40 lasers 334–6, 341–2 optical amplifiers 334–6, 340–1 rare-earth emissions and energy levels 333, 334 Xenon difluoride 386–8 Yttrium barium copper oxide 390 related structures and fluorine insertion 393–8 ZBLA(N) glasses 309, 310–11, 332–3 Zeolites, see Microporous materials