Electrochemistry of Functional Supramolecular Systems

ELECTROCHEMISTRY OF FUNCTIONAL SUPRAMOLECULAR SYSTEMS Edited by

Paola Ceroni Alberto Credi Margherita Venturi

The Wiley Series on Electrocatalysis and Electrochemistry

ELECTROCHEMISTRY OF FUNCTIONAL SUPRAMOLECULAR SYSTEMS

WILEY SERIES ON ELECTROCATALYSIS AND ELECTROCHEMISTRY Andrzej Wieckowski, Series Editor

Fuel Cell Catalysis: A Surface Science Approach, Edited by Marc T. M. Koper Electrochemistry of Functional Supramolecular Systems, Paola Ceroni, Alberto Credi, and Margherita Venturi Catalysis in Electrochemistry: From Fundamentals to Strategies for Fuel Cell Development, Elizabeth Santos and Wolfgang Schmickler Fuel Cell Science: Theory, Fundamentals, and Biocatalysis, Andrzej Wieckowski and Jens Norskov

ELECTROCHEMISTRY OF FUNCTIONAL SUPRAMOLECULAR SYSTEMS Edited by

Paola Ceroni Alberto Credi Margherita Venturi

The Wiley Series on Electrocatalysis and Electrochemistry

Copyright  2010 by John Wiley & Sons, Inc. All rights reserved. Published by John Wiley & Sons, Inc., Hoboken, New Jersey Published simultaneously in Canada 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, scanning, or otherwise, except as permitted under Section 107 or 108 of the 1976 United States Copyright Act, without either the prior written permission of the Publisher, or authorization through payment of the appropriate per-copy fee to the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, (978) 750-8400, fax (978) 750-4470, or on the web at www.copyright.com. Requests to the Publisher for permission should be addressed to the Permissions Department, John Wiley & Sons, Inc., 111 River Street, Hoboken, NJ 07030, (201) 748-6011, fax (201) 748-6008, or online at http://www.wiley.com/go/permissions. Limit of Liability/Disclaimer of Warranty: While the publisher and author have used their best efforts in preparing this book, they make no representations or warranties with respect to the accuracy or completeness of the contents of this book and specifically disclaim any implied warranties of merchantability or fitness for a particular purpose. No warranty may be created or extended by sales representatives or written sales materials. The advice and strategies contained herein may not be suitable for your situation. You should consult with a professional where appropriate. Neither the publisher nor author shall be liable for any loss of profit or any other commercial damages, including but not limited to special, incidental, consequential, or other damages. For general information on our other products and services, or technical support, please contact our Customer Care Department within the United States at (800) 762-2974, outside the United States at (317) 572-3993 or fax (317) 572-4002. Wiley also publishes its books in a variety of electronic formats. Some content that appears in print may not be available in electronic books. For more information about Wiley products, visit our web site at www.wiley.com Library of Congress Cataloging-in-Publication Data: ISBN 9780470255575 Printed in the United States of America 10 9 8

7 6 5 4

3 2 1

CONTENTS

Preface to the Wiley Series on Electrocatalysis and Electrochemistry

vii

Foreword

ix

Preface

xi

Contributors 1. Electrochemically Controlled H-Bonding

xiii 1

Diane K. Smith

2. Molecular Motions Driven by Transition Metal Redox Couples: Ion Translocation and Assembling– Disassembling of Dinuclear Double-Strand Helicates

33

Valeria Amendola and Luigi Fabbrizzi

3. Molecular Encapsulation of Redox-Active Guests

59

Angel E. Kaifer

4. Dendritic Encapsulation of Redox-Active Units

87

Christopher B. Gorman

5. Redox-Active Metal–Polypyridine Dendrimers as Light-Harvesting Antennae

121

Fausto Puntoriero, Scolastica Serroni, Francesco Nastasi, and Sebastiano Campagna

6. Dendrimers as Multielectron Storage Devices

145

Paola Ceroni and Margherita Venturi

7. Self-assembled Monolayers and Multilayers of Electroactive Thiols

185

Ibrahim Yildiz, Fran¸cisco M. Raymo and Massimiliano Lamberto

v

vi

CONTENTS

8. Electrochemistry of Carbon Nanoparticles

201

Luis Echegoyen, Amit Palkar, and Frederic Melin

9. Molecular Devices Based on Fullerenes and Carbon Nanotubes

229

Matteo Iurlo, Demis Paolucci, Massimo Marcaccio, and Francesco Paolucci

10. Functional Electroactive Biomolecules

261

Xiaomin Bin, Piotr Michal Diakowski, Kagan Kerman, Heinz-Bernhard Kraatz

11. Functional Nanoparticles as Catalysts and Sensors

301

Brian J. Jordan, Chandramouleeswaran Subramani, and Vincent M. Rotello

12. Biohybrid Electrochemical Devices

333

Ran Tel-Vered, Bilha Willner, and Itamar Willner

13. Electroactive Rotaxanes and Catenanes

377

Alberto Credi and Margherita Venturi

14. Electrochemically Driven Molecular Machines Based on Transition-metal Complexed Catenanes and Rotaxanes

425

Jean-Paul Collin, Fabien Durola, and Jean-Pierre Sauvage

15. Electroactive Molecules and Supramolecules for Information Processing and Storage

447

Guanxin Zhang, Deqing Zhang, and Daoben Zhu

16. Electrochemiluminescent Systems as Devices and Sensors

477

Andrzej Kapturkiewicz

17. Recent Developments in the Design of Dye-Sensitized Solar Cell Components

523

Stefano Caramori and Carlo Alberto Bignozzi

Index

581

PREFACE to the Wiley Series on Electrocatalysis and Electrochemistry

This series covers recent advances in electrocatalysis and electrochemistry and depicts prospects for their contribution to the present and future of the industrial world. It illustrates the transition of electrochemical sciences from a solid chapter of physical electrochemistry (covering mainly electron transfer reactions, concepts of electrode potentials, and structure of the electrical double layer) to the field in which electrochemical reactivity is shown as a unique chapter of heterogeneous catalysis, is supported by high-level theory, connects to other areas of science, and includes focus on electrode surface structure, reaction environment, and interfacial spectroscopy. The scope of this series ranges from electrocatalysis (practice, theory, relevance to fuel cell science and technology) to electrochemical charge transfer reactions, biocatalysis, and photoelectrochemistry. While individual volumes may look quite diverse, the series promises updated and overall synergistic reports on insights into further the understanding of properties of electrified solid/liquid systems. Readers of the series will also find strong reference to theoretical approaches for predicting electrocatalytic reactivity by such high-level theories as DFT. Beyond the theoretical perspective, further vehicles for growth are the sound experimental background and demonstration of significance of such topics as energy storage, syntheses of catalytic materials via rational design, nanometer-scale technologies, prospects in electrosynthesis, new instrumentation, surface modifications in basic research on charge transfer, and related interfacial reactivity. In this context, readers will notice that new methods that are being developed for a specific field may be readily adapted for application in others. Electrochemistry has benefited from numerous monographs and review articles due to its unique character and significance in the practical world (including electroanalysis). Electrocatalysis has also been the subject of individual reviews and compilations. The Wiley Series on Electrocatalysis and Electrochemistry is dedicated to the current activity by focusing each volume on a specific topic of choice. The chapters also demonstrate electrochemistry’s connections to other areas of chemistry and physics, such as biochemistry, chemical engineering, quantum mechanics, chemical physics, surface science, and biology, and illustrate the wide range of literature that each topic contains. While the title of each volume informs of the specific focus chosen by the volume editors and chapter authors, the integral outcome offers a broad-based analysis of the total development of the field. The progress of the series will provide a global definition of what electrocatalysis and electrochemistry are concerned with now and how they evolve with time. The purpose is manifold, vii

viii

PREFACE TO THE WILEY SERIES ON ELECTROCATALYSIS AND ELECTROCHEMISTRY

mainly to provide a modern reference for graduate instruction and for active researchers in the two disciplines, as well as to document that electrocatalysis and electrochemistry are dynamic fields that expand rapidly and likewise rapidly change in their scientific profiles. Creation of each volume required the editor’s involvement, vision, enthusiasm, and time. The Series Editor thanks all Volume Editors who graciously accepted his invitations. Special thanks are for Ms. Anita Lekhwani, the Series Acquisition Editor, who extended the invitation to the Series Editor and is a wonderful help in the Series assembling process. ANDRZEJ WIECKOWSKI Series Editor

FOREWORD

Like the currently popular area, called “nanoscience”, the field of “supramolecular chemistry” has rather hazy boundaries. Indeed, both areas now share much common ground in terms of the types of systems that are considered. From the beginning, electrochemistry, which provides a powerful complement to spectroscopic techniques, has played an important role in characterizing such systems and this very useful book goes considerably beyond the volume on this same topic by Kaifer and Go´mezKaifer that was published about 10 years ago. Some of the “classic” supramolecular chemistry topics such as rotaxanes, catenanes, host–guest interactions, dendrimers, and self-assembled monolayers remain, but now with important extensions into the realms of fullerenes, carbon nanotubes, and biomolecules, like DNA. These topics lead to considerations of supramolecular devices, for example for use as sensors, and to molecular machines. Not only is electrochemistry an excellent way of characterizing such systems, for example, via cyclic voltammetry, but in the world of molecular machines, it is also the most straightforward approach to providing the energy to power such devices. These topics then naturally lead to consideration of the conversion of electricity to light (electrochemiluminescence) and light to electricity (dye-sensitized solar cells) via electrochemical devices. While the latter are not fundamentally supramolecular systems, their design could certainly benefit from the considerations in the very detailed and authoritative treatments in this volume. The idea of integrated chemical systems, based on nanoscience and nanotechnology, was proposed a little over 20 years ago and was the subject of my 1994 monograph, but, so far, few such systems have reached widespread practical utilization. Nevertheless, the principles of such systems, for example for synthesis, analysis, and perhaps computation remain of interest, and supramolecular electrochemistry can play a major role in their development. I hope this important volume will go a long way toward introducing such principles to a wide audience, especially to the young people who are less burdened by impressions of what is impossible. The University of Texas at Austin

Allen J. Bard

ix

PREFACE

Supramolecular chemistry is a highly interdisciplinary field that has been developed at an astonishingly fast rate during the last three decades. In a historical perspective, as pointed out by Jean-Marie Lehn, supramolecular chemistry originated from Paul Ehrlich’s receptor idea, Alfred Werner’s coordination chemistry, and Emil Fischer’s lock-and-key image. It was only after 1970, however, that fundamental concepts such as molecular recognition, preorganization, self-assembly, and self-organization were introduced to chemistry; supramolecular chemistry then began to emerge as a welldefined discipline and was consecrated by the award of the Nobel Prize in Chemistry to Charles Pedersen, Donald Cram, and Jean-Marie Lehn in 1987. Supramolecular chemistry, according to its most popular definition, is “the chemistry beyond the molecule, bearing on organized entities of higher complexity that result from the association of two or more chemical species held together by intermolecular forces.” As the field developed, it became evident that a definition strictly based on the nature of the bond that links the components would be limiting. Many scientists, therefore, started to distinguish between what is molecular and what is supramolecular based on the degree of intercomponent interactions. In a general sense, one can say that with supramolecular chemistry there has been a shift in focus from molecules to molecular assemblies or multicomponent structures driven by the emergence of new functions. In the frame of research on supramolecular systems, the idea began to arise in a few laboratories that the concepts of “device” and “machine” could be applied at the molecular level. In other words, molecules might be used as building blocks for the assembly of multicomponent structures exhibiting novel and complex functions that arise from the cooperation of simpler functions performed by each component. This strategy, encouraged by a better understanding of biomolecular devices, has been implemented on a wide variety of chemical systems, leading to highly interesting results. As a matter of fact, the molecular bottom-up construction of nanoscale devices and machines has become one of the most stimulating challenges of nanoscience. Such achievements have been made possible because of the substantial progresses obtained in other areas of chemistry and physics—particularly concerning the synthesis and characterization of complex chemical systems, and the study of surfaces and interfaces. In this perspective, electrochemistry is a very powerful tool not only for characterizing a supramolecular system, but also for operating the device. Indeed, molecular devices, as their macroscopic counterparts, need energy to operate and signals to communicate with the operator. Electrochemistry can be an interesting xi

xii

PREFACE

answer to this dual requirement: it can be used to supply the energy needed to make the system work, and, by means of the various electrochemical techniques (e.g., voltammetry), it can also be used to “read” the state of the system, controlling and monitoring the operation performed by the device. Furthermore, electrodes represent one of the best ways to interface molecular-level systems to the macroscopic world, a feature that is important for future applications. Hence, it is not surprising that the marriage of electrochemistry and supramolecular chemistry has produced a wealth of very interesting devices and functions, thereby generating new scientific knowledge and raising expectations for practical applications in energy conversion, information and communication technologies, advanced materials, diagnostics, and medicine. Our aim with this book is to provide the reader with an overview of current electrochemical research applied to multicomponent chemical systems, with particular attention to properties and functions, and to strengthen the contacts between the electrochemical community and the researchers engaged in the field of nanoscience. Although the text covers a wide range of topics with contributions from leading authorities in their respective fields, it does not even attempt to be a comprehensive book on supramolecular electrochemistry. Rather, we would like to give the reader a flavor of the level of creativity and ingenuity reached by scientists working in this area. We hope that the book will be useful as a reference not only for experienced researchers, but also for graduate students and postdoctoral fellows who are interested in exploring electrochemistry at its frontiers with supramolecular chemistry, materials science, and biochemistry. It may also be a useful complement for students attending nanoscience and nanotechnology courses. The 17 chapters of the book are not grouped in sections but they are somehow logically ordered. The initial contributions, describing basic science investigations on systems in solution, are followed by chapters dealing with less conventional multicomponent architectures and/or environments. The final part contains contributions on devices and systems of high complexity and/or applicative interest. Although the book does not include introductory or tutorial sections, most chapters begin with a discussion of the basic concepts that are relevant for the presented topics. We hope that these sections will make the book comprehensible also to nonspecialists. We would like to express our gratitude to the distinguished colleagues and friends who contributed the chapters: their commitment is indeed a fundamental ingredient of this initiative. We also thank people at Wiley for their assistance during the various phases of the editorial work. Finally, we would like to thank our families because their love and patience are an invaluable support for our professional activity. Bologna, June 2009

PAOLA CERONI ALBERTO CREDI MARGHERITA VENTURI

CONTRIBUTORS

Valeria Amendola, Dipartimento di Chimica Generale, Universita di Pavia, Pavia, Italy Carlo Alberto Bignozzi, Dipartimento di Chimica, Universita di Ferrara, Ferrara, Italy Xiaomin Bin, Department of Chemistry, The University of Western Ontario, London, Ontario, Canada Sebastiano Campagna, Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Universit a di Messina, Messina, Italy Stefano Caramori, Dipartimento di Chimica, Universita di Ferrara, Ferrara, Italy Paola Ceroni, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum Universit a di Bologna, Bologna, Italy Jean-Paul Collin, Laboratoire de Chimie Organo-Minerale, UMR 7177 du CNRS, Faculte de Chimie, Universite de Strasbourg, Strasbourg Cedex, France Alberto Credi, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum, Universit a di Bologna, Bologna, Italy Piotr Michal Diakowski, Department of Chemistry, The University of Western Ontario, London, Ontario, Canada Fabien Durola, Laboratoire de Chimie Organo-Minerale, UMR 7177 du CNRS, Faculte de Chimie, Universite de Strasbourg, Strasbourg Cedex, France Luis Echegoyen, Department of Chemistry, Clemson University, Clemson, SC, USA Luigi Fabbrizzi, Dipartimento di Chimica Generale, Universita di Pavia, Pavia, Italy Christopher B. Gorman, Department of Chemistry, North Carolina State University, Raleigh, NC, USA Matteo Iurlo, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum, Universit a di Bologna, Bologna, Italy Brian J. Jordan, Department of Chemistry, University of Massachusetts, Amherst, MA, USA xiii

xiv

CONTRIBUTORS

Angel E. Kaifer, Center for Supramolecular Science, Department of Chemistry, University of Miami, Coral Gables, FL, USA Andrzej Kapturkiewicz, Institute of Physical Chemistry, Polish Academy of Sciences, Warsaw, Poland and Institute of Chemistry, University of Podlasie, Siedlce, Poland Kagan Kerman, Department of Chemistry, The University of Western Ontario, London, Ontario, Canada Heinz-Bernhard Kraatz, Department of Chemistry, The University of Western Ontario, London, Ontario, Canada Massimiliano Lamberto, Department of Chemistry, Medical Technology and Physics, Monmouth University, West Long Branch, NJ, USA Massimo Marcaccio, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum, Universit a di Bologna, Bologna, Italy Frederic Melin, Department of Chemistry, Clemson University, Clemson, SC, USA Francesco Nastasi, Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Universit a di Messina, Messina, Italy Amit Palkar, Department of Chemistry, Clemson University, Clemson, SC, USA Demis Paolucci, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum, Universit a di Bologna, Bologna, Italy Francesco Paolucci, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum, Universit a di Bologna, Bologna, Italy Fausto Puntoriero, Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Universit a di Messina, Messina, Italy Fran¸cisco M. Raymo, Department of Chemistry, University of Miami, Coral Gables, FL, USA Vincent M. Rotello, Department of Chemistry, University of Massachusetts, Amherst, MA, USA Jean-Pierre Sauvage, Laboratoire de Chimie Organo-Minerale, UMR 7177 du CNRS, Faculte de Chimie, Universite de Strasbourg, Strasbourg Cedex, France Scolastica Serroni, Dipartimento di Chimica Inorganica, Chimica Analitica e Chimica Fisica, Universit a di Messina, Messina, Italy Diane K. Smith, Department of Chemistry and Biochemistry, San Diego State University, San Diego, CA, USA Chandramouleeswaran Subramani, Department of Chemistry, University of Massachusetts, Amherst, MA, USA

CONTRIBUTORS

xv

Margherita Venturi, Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum, Universit a di Bologna, Bologna, Italy Ran Tel-Vered, Institute of Chemistry, The Hebrew University of Jerusalem, Jerusalem, Israel Bilha Willner, Institute of Chemistry, The Hebrew University of Jerusalem, Jerusalem, Israel Itamar Willner, Institute of Chemistry, The Hebrew University of Jerusalem, Jerusalem, Israel Ibrahim Yildiz, Department of Chemistry, University of Miami, Coral Gables, FL, USA Deqing Zhang, Beijing National Laboratory for Molecular Sciences, Organic Solids Laboratory, Institute of Chemistry, Chinese Academy of Sciences, Beijing, China Guanxin Zhang, Beijing National Laboratory for Molecular Sciences, Organic Solids Laboratory, Institute of Chemistry, Chinese Academy of Sciences, Beijing, China Daoben Zhu, Beijing National Laboratory for Molecular Sciences, Organic Solids Laboratory, Institute of Chemistry, Chinese Academy of Sciences, Beijing, China

CHAPTER 1

Electrochemically Controlled H-Bonding DIANE K. SMITH Department of Chemistry and Biochemistry, San Diego State University, San Diego, CA, USA

1.1

INTRODUCTION

Due to their strength and directionality, hydrogen bonds are one of the most important and useful types of intermolecular interactions available for the construction of supramolecular complexes. The iconic examples of DNA base pairing and the formation of secondary structure in proteins provide ample proof of their utility for the assembly of well-defined, functional structures. Examples1 of the use of hydrogen bonds in synthetic, solution-phase supramolecular chemistry range from H-bonded dimers2 held together by up to 6 H-bonds3 to large “rosette” assemblies constructed from up to 15 components and 72 H-bonds.4 A wide variety of open H-bonded structures have also been prepared, including those that self-assemble into capsules of various sizes and shapes5,6 and cyclic peptides that assemble into hollow tubes.7 Although, from a purely chemical point of view, learning how to create these complicated supramolecular structures has its own value, there are plenty of more practical reasons to investigate this chemistry. In the short term, these include catalysis and sensor applications, and in the long term, molecular electronics and molecular machines. With perhaps the exception of catalysis, all these applications will require some sort of signal transduction to allow for communication with the supramolecular device. This, of course, is one of the main reasons that electrochemistry is useful for supramolecular chemistry. Electron transfer provides a wellunderstood and very sensitive method to both communicate with supramolecular assemblies and control their structure.8 However, although electrochemistry can be used for the above, it will do so only if this functionality is designed into the structure. At minimum, two requirements Electrochemistry of Functional Supramolecular Systems. Edited by Paola Ceroni, Alberto Credi, and Margherita Venturi Copyright  2010 John Wiley & Sons, Inc.

1

2

ELECTROCHEMICALLY CONTROLLED H-BONDING

must be met. First, a reversible redox couple must be present as part of the structure. Reversible in this context means that both oxidation states are chemically stable under the experimental conditions used and that the electron transfer kinetics are reasonably fast. Second, reduction or oxidation of the redox couple must significantly perturb the strength of important binding interactions holding the assembly together. The most straightforward type of interaction to perturb electrochemically are ion–ion interactions, but the electrostatic nature of a hydrogen bond makes it a close second, while also allowing for neutral molecules as binding partners. In this chapter, the basic principles behind electrochemically controlled H-bonding will first be described, along with some simple, illustrative examples and a brief discussion of the use of cyclic voltammetry to characterize such systems. Next, some general considerations regarding the design of these systems are discussed: the properties of the redox couple, the structures of host and guest, the choice of solvent and electrolyte, and the possibility for proton transfer. Finally, a selection of electrochemically controlled H-bonding systems will be described, organized by the nature of the binding partner and the type of redox couple.

1.2

BASIC PRINCIPLES

A H-bond is a favorable interaction formed between a relatively positively charged hydrogen atom in a polar bond, XH, and either a lone pair on a relatively negatively charged atom, Y (Eq. 1.1), or a highly polarizable pi-bonding electron pair (Eq. 1.2).9 As this description implies, there is a high degree of electrostatic character associated with H-bonding, although only the very weakest H-bonds are purely electrostatic. Strong H-bonds range in strength from 15 to 40 kcal/mol and are considered to be mainly covalent in character. These bonds are characterized by close to linear XHY bond angles and XY distances that are substantially smaller than the sum of the van der Waals radii. These are typically formed from ionic species, for example, N þ H or O þ H as the H-donor and/or F, O, or N as the H-accepting atom. Moderate strength H-bonds range in strength from 4 to 15 kcal/mol and are mostly electrostatic in character. They can show a greater range in bond angles, from 130 to 180 , and will have bond lengths that are slightly smaller than the sum of the van der Waals radii. Typical examples are H-bonds formed from neutral oxygen and nitrogen functional groups, for example, NH and OH as the donor group and uncharged N and O as the accepting atoms. Weak H-bonds, those less than 4 kcal/mol in strength, are almost purely electrostatic in character and are characterized by a range of bond angles and XY distances that may be greater than the sum of the van der Waals radii. These are formed from the weaker H-donors such as CH and/or weaker acceptors such as pi electron pairs. +

X H

+

+

Y R

X H

Y R

ð1:1Þ

1.2

+

+

X H

1.2.1

+

+

R Y Y R

X H

3

BASIC PRINCIPLES

R Y Y R

ð1:2Þ

Direct Perturbation of Hydrogen Bonds

Although, from the above discussion, H-bonds are generally not purely electrostatic, a favorable electrostatic interaction plays an important role in all types of H-bonds. This provides a very simple way to think about how to use electron transfer to directly perturb the strength of H-bonds. There are two main ways to do this as shown in Scheme 1.1. The first is to make a H-acceptor a better acceptor by using a reduction reaction to increase the negative charge on the H-accepting atom (Scheme 1.1a). The second is to make the H-donor a better donor by using oxidation to increase the positive charge on a H-donating functional group (Scheme 1.1b). Alternatively, electron transfer can be used to weaken H-bonds through the opposite effects, making a H-donor a weaker donor through reduction or making a H-acceptor a weaker acceptor through oxidation. A simple example of reduction-based, electrochemically controlled H-bonding is provided by nitrobenzenes.10 Despite the standard Lewis structure that places a formal negative charge on one of the oxygens and a positive charge on the nitrogen, nitro groups are generally weak H-acceptors in solution because the NO bond is not very polar. However, reduction of an aromatic nitro compound to its radical anion greatly increases the negative charge on the oxygens, resulting in much stronger H-bonding to a H-donating guest. In the case of nitroaniline, 1 (X ¼ NH2), with 1,3-diphenylurea, 2 (Eq. 1.3), the equilibrium constant for H-bonding in 0.1 M NBu4PF6/DMF goes from 10 times the host concentration. If these conditions are not met, then an alternative strategy is needed. The most powerful is to use CV simulation software to fit the experimental CVs to the square scheme or a more complicated mechanism if necessary. This method allows determination of the thermodynamic parameters and possibly the kinetic parameters as well. An example of the use of CV simulation to determine binding constants and explain more complicated CV behavior is shown in Fig. 1.2. These are CVs observed

Figure 1.2 Experimental (lines) and simulated (dots) CVs of 1 mM 3 in 0.1 M NBu4B(C6F5)4/ CH2Cl2 in the presence of different amounts of 4: (a) 0 mM, (b) 0.5 mM, (c) 1 mM, and (d) 100 mM. 500 mV/s scan rate.11

8

ELECTROCHEMICALLY CONTROLLED H-BONDING

for oxidation of the electroactive urea, 3, discussed previously in the presence of the cyclic diamide, 4.11 Now the guest is expected to bind more strongly to the oxidized form (Eq. 1.4), resulting in a negative shift in E1/2 upon addition of guest, the opposite of the effect seen in Fig. 1.1. Note that under the experimental conditions used, the CV wave for the urea by itself, scan (a), is broad. With half equivalent of guest, scan (b), an even broader CV results, with a new shoulder negative of the original wave. With continued addition of guest, the wave now sharpens up and continues to move negative. In order to explain the original broad wave observed for the urea, it has been proposed that formation of dimer 5 between an unoxidized urea and an oxidized urea is quite favorable. This makes it easier to oxidize the first half of the urea (since the dimer will result) than the second half of the urea (which requires breakup of the dimer). The result is an unusually broad CV wave. Addition of greater than 1 equivalent of the guest breaks up the dimer resulting in a sharper wave. The viability of this mechanism is confirmed by showing that CVs simulated using this mechanism (dots in Fig. 1.2) give good fits to the experimental CVs (lines) with reasonable values for the different reaction parameters.

H

Ph N

Me2NPh

N

O

H

H

O 5

N

Ph H N PhNMe2

While cyclic voltammetry is clearly the most important technique for studying electrochemically controlled H-bonding, other physical methods, in particular spectroscopic techniques, can be quite helpful. These can be used to provide structural information on the host–guest complex, and also to provide another means to determine binding constants in at least one of the redox states. For H-bonded supramolecular complexes, the most commonly used technique is simply 1 H NMR, since the chemical shifts of hydrogens involved in the H-bonding will be very sensitive to the presence of the binding partner. However, a significant limitation for the use of NMR in electrochemical H-bonding studies is that it can only be used with diamagnetic systems, and in many cases the stronger binding state is paramagnetic. 1.4 GENERAL CONSIDERATIONS FOR THE DESIGN OF ELECTROCHEMICALLY CONTROLLED HYDROGEN-BONDED ASSEMBLIES 1.4.1

The Redox Couple

First and foremost, the redox couple must be reversible, meaning relatively fast electron transfer reactions and products that are stable in both oxidation states. As discussed previously, for systems in which the hydrogen bonds are to be directly perturbed, the redox reaction must affect the charge distribution on the atoms involved

1.4

GENERAL CONSIDERATIONS FOR THE DESIGN

9

TABLE 1.1 Shift in Half-Wave Potential, DE1/2, for Different para-Substituted Nitrobenzenes in the Presence of 1,3-Diphenylureaa Substituent NH2 CH3O CH3 H CF3

E1/2 (V) versus Fc

DE1/2 (mV)

1.870 1.698 1.635 1.582 1.362

197 164 156 153 93

a 1 mM nitrobenzene in 0.1 M NBu4PF6/DMF þ 50 mM 1,3-diphenylurea. Values are the averages of at least three independent measurements.

in one or more of the hydrogen bonds. This can be done through what can be described as either an inductive or a resonance effect. The nitrobenzene (Eq. 1.3)10 and dimethylaminophenylurea systems (Eq. 1.4)11 that were discussed in the previous sections are good examples of resonance effects. In these cases, the electron transfer directly places charge on the atoms involved in H-bonding, which are an integral part of the redox couple. One interesting possibility is to couple this with simple substituent effects in order to further “tune” the binding strength. This hypothesis has only been explored in two systems so far, with mixed results. One is the nitrobenzene system, where the strategy does appear to work, as shown by the data in Table 1.1.10 This table gives the observed E1/2 of the nitrobenzene 0/1 redox couple, along with the observed DE1/2 in the presence of 50 equivalents of 1,3-diphenylurea for five different p-substituted nitrobenzenes. As expected, the E1/2 values fall in order of the electron-donating/withdrawing strength of the substituents, with the nitrobenzene with the most electron-donating substituent, NH2, being the hardest to reduce and that with the most electron-withdrawing substituent, CF3, being the easiest to reduce. Interestingly, the DE1/2 values fall in the same order, with the NH2 derivative giving the largest shift and the CF3 the smallest. This is consistent with the electron-donating NH2 forcing more of the negative charge in the radical anion onto the oxygens leading to an increase in H-bond strength, and with the electron-withdrawing CF3 removing some of the negative charge from the oxygens leading to a decrease in H-bond strength. However, although tuning DE1/2 values with substituents appears to be quite effective in the nitrobenzene system, an even more in-depth substituent study for the flavin/amidopyridine H-bonding system,15 which will be discussed in detail later in the chapter, shows the expected E1/2 dependence, but no clear trend in the DE1/2 values. It therefore remains to be seen how generally useful this strategy may be. It is also possible to affect H-bond strength electrochemically without the H-bonding site being an integral part of the redox couple. Good examples of this are found in the many redox-dependent receptors that utilize metallocenes, primarily ferrocene and cobaltocenium, as the redox couple. These are primarily used in ion receptors, but examples of metallocene receptors that show a significant redox dependence with neutral guests are the cobaltocenium and ferrocene diamides,

10

ELECTROCHEMICALLY CONTROLLED H-BONDING

6 and 7. Both metallocenes undergo a reversible one-electron transfer reaction between 0 and þ 1 charge states. However, since cobalt has one more valence electron than iron, the cobalt derivative has the optimum 18 valence electron configuration and is therefore most stable in the þ 1 state, whereas the iron analog has the optimum 18 valence electron configuration and is most stable in the zero charge state. O N H O 6 M = Co+ 7 M = Fe

N H O (CH2)3

M O H N O

O H N

8

Tucker and coworkers have shown that both of these compounds bind glutaric acid, 8, through H-bonding between the carboxylic acids and the diamidopyridine.161 H NMR titrations in CDCl3/DMSO (0.5%) indicate that the cobaltocenium host binds 20 times more strongly than the ferrocene host (Kox of 6 ¼ 9.8  104 M1; Kred of 7 ¼ 4.6  103 M1). This is consistent with the positive charge on the cobaltocenium, which would be expected to further polarize the NH amide bond through an electron-withdrawing effect, resulting in stronger H-bonding to the carbonyl oxygen of the acid. In fact, crystal structures show that the amide H-bonds are shorter in the cobaltocenium derivative than the ferrocene, whereas the pyridine H-bonds are about the same. Reduction of the cobaltocenium to the zero state, and oxidation of the ferrocene to the þ 1 state would be expected to switch the binding preferences, and indeed the CV studies in CH2Cl2/DMSO (0.5%) indicate that this is the case. Addition of excess glutaric acid to both compounds results in the same maximum DE1/2 of 90 mV. This corresponds to a 30-fold difference in binding strength between oxidation states, with the þ 1 state being the stronger binding one in both cases. This indicates that the difference in binding between oxidation states for this system is simply due to the change in charge. Unlike the nitrobenzene example, it is independent of variation in structure (Co versus Fc) or actual E1/2 since the E1/2 of Co(Cp)2 þ 1/0 is quite negative of the E1/2 of Fe(Cp)2 þ 1/0. In comparing the two different ways the redox couple can perturb H-bonding strength, there are clearly pros and cons to both. The resonance approach, in which the H-bonding sites are an integral part of the redox couple, is likely to give a stronger redox perturbation than the inductive approach, where the redox couple is simply in conjugation with the H-bonding site. The existing examples bear this out. With neutral guests, where H-bonding is the main binding interaction, the electroactive hosts that have been shown to give the very large DE1/2 values are all those in which the H-bonding site is part of the redox couple. The disadvantage is that this really

1.4

GENERAL CONSIDERATIONS FOR THE DESIGN

11

puts severe limitations on host/guest combinations possible. The advantage of the inductive approach is that the binding site can be designed somewhat independent of the redox couple, making it, in principle, easier to tailor the host for different guests. The other advantage is that very inert redox couples such as the metallocenes can be used, so there is less chance of unwanted reactions taking place. Another factor to take into consideration when choosing a redox couple is the availability of additional electron transfer equilibria. The redox-dependent hosts considered so far, nitrobenzenes, 1, the dimethylaminophenylurea, 3, and the metallocene diamides, 6 and 7, just undergo one reversible electron transfer reaction. However, many organic redox couples undergo at least two successive electron transfers. The prototypical example are the quinones, 9, which in aprotic solvents undergo two reversible reductions, first to the radical anion and then to dianion (Eq. 1.8). p-Phenylenediamines, 10, provide a related example of a redox couple that undergoes two reversible oxidations, giving first a radical cation and then a dication (Eq. 1.9). (Note that the dimethylaminophenylurea, 3, is a phenylenediamine derivative, but the second oxidation is irreversible in this case.) In general, the more highly charged dianions and dications will be significantly more basic or acidic than the mono-ions, and therefore will H-bond more strongly to a given guest, resulting in much larger DE1/2 values in the 1/2 or þ 1/ þ 2 redox couple than the 0/1 or 0/ þ 1 redox couples. This phenomena has been particularly well studied with quinones.17 Addition of weak H-donors, such as alcohols, produces large positive shifts in the 1/2 redox potential with little change in the 0/1 potential. Similar results are observed with p-phenylenediamines and weak H-acceptors such as tertiary amides.18 O

O +

e–

O +

e–

ð1:7Þ O

O

O

NH2

NH2

9 NH2 – e–

– e–

ð1:8Þ NH2

NH2

NH2

10

Based on the above discussion, there is an advantage in using redox couples with multiple electron transfer equilibria since this makes incremental control of binding strength possible. However, caution must also be used, since the stronger H-bonding also means that the response will be much less selective, in that weaker H-donors/ acceptors could also cause a significant shift. In addition, because the dianions and dications will be more reactive, there is a greater likelihood of proton transfer and other chemical reactions taking place at the second electron transfer, which could

12

ELECTROCHEMICALLY CONTROLLED H-BONDING

compromise the reversibility of the system. These issues are problematic for practical sensor applications, where the redox response should be highly selective for a particular analyte in a range of different environments. However, this may not be as much an issue in other applications, where the environment can be controlled to prevent unwanted interactions with the more reactive oxidation states. 1.4.2

Host/Guest Structure

In order to achieve strong, selective binding with electroactive hosts, just as in other types of molecular recognition, there needs to be a high degree of complementarity between the host and the guest. This means that the binding site needs to have the proper shape, along with the correct positioning of functionality to provide multiple favorable contacts with the target guest. The principle of preorganization is also important, which says that binding strength will be maximized when the host is “preorganized” to fit the guest because this will minimize the entropy loss upon formation of the host–guest complex. For this reason, rigidity in both the host and if possible the guest is generally desirable. Of course, in order to produce a redox signal for binding, at least one of the binding interactions must be strongly perturbed by a reversible electron transfer reaction. As discussed in the previous section, a greater effect will generally be observed at the second electron transfer (if available), producing a larger DE1/2 and making the redox potential of the host sensitive to the guest over a larger range of concentrations. However, this will also likely make the response much less selective, so for sensor applications, the first electron transfer will probably need to be used. This means that a big change in binding strength is needed based on what will typically be a 0/1 or 0/ þ 1 redox couple. In looking at existing electrochemically controlled H-bonding systems in which significant E1/2 shifts have been observed at the first electron transfer, two minimum features appear necessary. First, there needs to be the possibility of at least two strong, almost linear H-bonds formed between host and guest, and, second, the electron transfer reaction needs to strongly perturb at least one and preferably two H-bonds between host and guest. Examination of the systems discussed so far in this chapter, as well as those yet to be discussed, shows that all obey these basic design principles. A nice demonstration of the above criteria is provided by a study of the redoxdependent H-bonding properties of 9,10-phenanthrenequinone, 11, and related compounds.19 As with nitrobenzenes, diarylureas such as 2 make excellent binding partners, providing appropriately positioned amide NHs to be able to simultaneously H-bond with both carbonyl oxygens. Also like nitrobenzene, reduction of phenanthrenequinone produces a radical anion, with enhanced negative charge on the oxygens. This should result in stronger H-bonding in the radical anion state (Eq. 1.10). Consistent with this, addition of 10 equivalents of 1,3-diphenylurea causes the E1/2 of the phenanthrenequinone 0/1 couple to shift positive by 200 mV in CH2Cl2. Significant positive E1/2 shifts are also observed in the more polar DMF. CV simulation of the latter gives Kox ¼ 1 M1 and Kred ¼ 905 M1 in 0.1 M NBu4PF6/DMF.

1.4

13

GENERAL CONSIDERATIONS FOR THE DESIGN

Ph O

O

+ e–

+2

O

H N

O

H N

O O

O

2 Ph 11

ð1:10Þ In order to investigate the structural requirements for the strong redox dependence, DE1/2 values were also measured for anthraquinone, 12, and benzyl, 13, in the presence of 5 equivalents of diphenylurea in DMF. Under these conditions, phenanthrenequinone gives a shift of 61 mV, whereas anthraquinone gives a shift of only 8 mVand benzil 5 mV. Unlike phenanthrenequinone, the urea can only H-bond to one carbonyl oxygen at a time with anthraquinone. Two bifurcated H-bonds are possible, but these together would be much weaker than the two close to linear H-bonds possible with o-quinones. A similar situation arises with benzil, since rotation about the central CC bond will be hindered in the radical anion and the favored conformation will have the oxygens trans due to electrostatic repulsion and steric effects. O O

O O 12

13

Another interesting comparison is the difference in DE1/2 observed for phenanthrenequinone with the two pyridylureas 14 and 15. The electronic character of both compounds should be similar, but with 14 a strong intramolecular H-bond will be formed between the 2-pyridyl N and one of the urea NHs. This means that only one NH will be available for H-bonding to the phenanthrenequinone radical anion. This has a huge effect on the DE1/2 values, with 5 equivalents of 14 producing only a 16 mV shift in DMF compared to an 85 mV shift with 15. The latter is actually larger than 1,3diphenylurea. This can be explained by the greater electronegativity of N compared to C, making the pyridyl group more electron withdrawing than a phenyl group. O H

N N N

N H

N H 14

1.4.3

O N H 15

Solvent and Electrolyte

In addition to the structures of host and guest, another important consideration for studies of electrochemically controlled H-bonding is the solvent system.

14

ELECTROCHEMICALLY CONTROLLED H-BONDING

Binding constants for supramolecular assemblies are always strongly solvent dependent because binding of the host to the guest will be in competition with solvation. Since H-bonding has a strong electrostatic component, more polar solvents will solvate the binding sites more effectively than less polar solvents, thereby decreasing the equilibrium constant for host–guest binding, whereas less polar solvents will tend to increase the binding strength. For nonelectrochemical studies, chloroform (more specifically, deuterated chloroform for 1 H NMR studies) seems to be the most common solvent. For electrochemical studies, a less resistive solvent is needed, with dichloromethane being the most common choice. Other aprotic solvents such as acetonitrile or dimethylformamide are also used. Sometimes, these can help to simplify the behavior by decreasing the amount of host–host or guest–guest interaction that occurs in the less polar dichloromethane. Another important consideration for electrochemical studies is the choice of electrolyte, the ionic compound that is added to maintain electroneutrality and provide a means of charge flow through solution. Because at least one of the oxidation states of the host will be charged, it is possible that ion–ion interactions will play a role in the observed electrochemistry. Indeed, this is useful if the objective is to design a redox-dependent ion receptor, but it can be an interference if the guest is neutral. The most common electrolyte used today for electrochemical studies in aprotic solvents is tetrabutylammonium hexafluorophosphate, NBu4PF6. The use of NBu4 þ is a good choice for reduction-based systems since it is a very large cation and is unlikely to interact significantly with the anionic form of the host. For oxidation-based systems, which generally involved cationic intermediates, it is the anion in the electrolyte that is of issue. PF6 has generally been considered to also be a large and noncoordinating ion. However, recent evidence suggests that it is not nearly as inert as once believed. In particular, electrochemical studies with organometallic cations show that PF6 and other “large” anions such as ClO4 and BF4 interact with þ 2 and þ 3 charged species enough to have a large effect on the observed E1/2 values, sometimes also causing distortion in the CV wave shapes and/or precipitation of the cationic products onto the electrode.20,21 If good H-donor groups are part of the redox couple, it might be expected that even þ 1 charge species could show strong interactions with these anions. This has recently been shown to be the case with the electroactive urea 3 discussed earlier (Eq. 1.4).11 In the results described previously (Fig. 1.2), the electrolyte was 0.1 M NBu4B(C6F5)4/ CH2Cl2. B(C6F5)4 is a very large anion, much larger than the more commonly used PF6. If the same experiments are run with NBu4PF6 or NBu4ClO4, much smaller E1/2 shifts are observed with the same guest. For example, addition of 1 equivalent of guest 4 to 3 gives DE1/2 values of 2 and 16 mV using NBu4ClO4 or NBu4PF6, respectively, but the same amount of guest gives 109 mV shift when NBu4B (C6F5)4 is used as the electrolyte. These results indicate that the smaller anions (which are present at much larger concentrations than the guest) interact strongly enough with the oxidized urea to block interactions with the guest. The much larger anion B(C6F5)4 does not do this, with the result that significantly larger shifts are observed in the presence of this anion. Based on these results, it would seem prudent

1.4

GENERAL CONSIDERATIONS FOR THE DESIGN

15

to at least test oxidation-based systems with these very large anion electrolytes, particularly if a less polar solvent such as CH2Cl2 is being used. 1.4.4

H-Bonding Versus Proton Transfer

Although it seems to be often pointed out that H-bonding and proton transfer are not the same thing (and obviously they are not), there is clearly a close relationship between the two, with stronger acids generally being better H-donors and stronger bases being better H-acceptors. As the pKa of a H-donor guest decreases, the E1/2 shift between the guest and the host will increase but so will the possibility of proton transfer. For organic 0/1 or 0/ þ 1 redox couples, the occurrence of proton transfer is usually quite obvious. Instead of seeing simply a shift in potential of the reversible, one-electron CV wave upon addition of the guest, the wave doubles in size and becomes irreversible. This behavior, which was originally observed upon addition of acids to aromatic anions, is due to an ECEC mechanism. (E stands for electron transfer reaction and C stands for a chemical reaction such as proton transfer.) In the case of reductions, the first step is electron transfer to form a radical anion (Eq. 1.11). Normally, addition of a second electron occurs at a more negative potential because it will be harder to add an electron to an already negatively charged species. However, protonation of the radical (Eq. 1.12) gives an uncharged radical that is typically easier to reduce than the starting species. The result is immediate addition of a second electron (Eq. 1.13) resulting in a two-electron CV wave. Since an acid strong enough to protonate the radical is usually strong enough to protonate the anion product, the final product will generally be the two-electron, two-proton reduced species, RH2 (Eq. 1.14). The reduction is irreversible because it will be much harder to oxidize RH2 than R. A similar mechanism can be observed with oxidations in the presence of bases capable of removing a proton from the radical cation intermediate. R þ e  ! R

ð1:11Þ

R þ H þ ! RH

ð1:12Þ

RH þ e ! RH

ð1:13Þ

RH þ H þ ! RH2

ð1:14Þ

Since the products of the second electron transfers are typically more basic or more acidic than those of the first electron transfer, it is possible to add a guest that is only capable of proton transfer with the product of the second electron transfer. It is important to note that in this case the result of proton transfer may be very similar to that of H-bonding, with simply a positive E1/2 shift for the second wave in a reduction or a negative shift for an oxidation. In this case, information about pKa values is very helpful for trying to sort out whether it is H-bonding or proton transfer that causes the shift.

16

ELECTROCHEMICALLY CONTROLLED H-BONDING

1.5 EXAMPLES OF ELECTROCHEMICALLY CONTROLLED H-BONDING SYSTEMS WITH ANIONIC GUESTS Although the examples of electrochemically controlled H-bonding discussed so far in this chapter have been with neutral guests, the reality is that the majority of the work that has been done in this area has been with ionic, and in particular anionic guests. This work started in the late 1980s, several years before the first reports with neutral guests. It is still an incredibly active area, with numerous new receptors reported each year. Work through the 1990s has been covered in review articles.22,23 The first reported example of anion binding with a synthetic organic macrocycle was the protonated cryptand 16, reported by Park and Simmons in 1968.24 1 HNMR evidence was presented showing that this compound binds chloride by a combination of H-bonding and favorable electrostatic interaction. Although numerous other anion receptors have been reported since then, including those that utilize Lewis acid and hydrophobic interactions, the two binding interactions seen in this first example, ion–ion and H-bonding, have turned out to be the mainstays of anion recognition. This leads naturally to electrochemically controlled H-bonding since adding an electroactive group provides a simple way to perturb both electrostatic and H-bonding interactions. O (CH2)9 N H

Cl (CH2)9 (CH2)9 16

1.5.1

O O

O

Co+

Co+

H N

O

O

O

O 17

Cobaltocenium

The first redox-active receptor for anions was the bis-cobaltocenium macrocylic ester 17 reported by Beer and Keefe in 1989.25 FAB-MS and FT-IR evidence indicates that 17 can bind anions in the oxidized state. Reduction of the cobaltoceniums to the zero state would be expected to weaken this interaction, and indeed a modest 45 mV E1/2 shift is observed in acetonitrile upon addition of 4 equivalents of Br. In the above example, binding and redox dependence rely solely on perturbation of electrostatic interactions. A few years later, in 1992, Beer and coworkers reported similar or better redox dependence with halides using simpler acyclic cobaltoceniums.26,27 The key was the addition of an amide NH, providing a H-bonding site. For example, the simple diamide cobaltocene 18 gives a 60 mV shift in E1/2 of the cobaltocenium upon addition of 4 equivalents of Br in acetonitrile, corresponding to at least a 10-fold decrease in binding strength. The analogous NMe compound results in a smaller than 5 mV shift. A later crystal structure of 18 with Br clearly

1.5 EXAMPLES OF ELECTROCHEMICALLY CONTROLLED H-BONDING SYSTEMS

17

shows the H-bond with the amide NH.28 In addition, a significantly larger shift of 240 mV is observed with 18 and H2PO4, due to the greater H-bonding capability of this anion compared to simple halides.

O

O N H H N

Co+

18

1.5.2

Fe

19

N H H N

N H

H N

O

O

Ferrocene

The cobaltocenium anion receptors are examples of electrochemically controlled H-bonding systems in which the electron transfer decreases H-bonding strength. By switching to ferrocene as the redox-active group, anion receptors can be made in which electron transfer increases binding strength. This has proven to be a particular popular strategy and there are now numerous examples of ferrocene-based, redoxdependent anion receptors in the literature. A generally successful strategy, learned from the early cobaltocenium examples, is to simply attach good hydrogen donors, such as amides, secondary amines, guanidiniums, and so on, to one or both cyclopentadienes. A representative example is the diurea-substituted ferrocene, 19, reported by Molina and coworkers,29 which shows a significant redox-dependent response to F (DE1/2 ¼ 208 mV) and H2PO4 (DE1/2 ¼ 90 mV) in DMSO, but not to Cl, Br, AcO, NO3, and HSO4. This is an example of a receptor where strong binding is observed in both oxidation states, resulting in two CV waves being observed at less than 1 equivalent of guest. Most redox-dependent receptors have structural features similar to 17–19, which are either a macrocyclic or a cleft-type binding site with one or two redox-active groups directly attached. An interesting alternative, introduced by Astruc and coworkers, are the ferrocene-terminated dendrimers, such as 20.30,31 Although each dendrimer contains multiple ferrocenes, just one CV oxidation wave is generally observed, indicating that there is no significant interaction between the ferrocenes. Addition of a good H-bonding anion either causes this wave to shift negative or produces a new CV wave. One of the interesting results is that there is generally a “positive dendritic effect,” in that higher dendrimer generations show a larger maximum shift. For example, Fig. 1.3 shows the effect of different generations of amine dendrimers on the observed DE1/2 caused by addition of HSO4. (18-Fc is 20, which has 18 ferrocenes, 9-Fc is the previous generation with 9 ferrocenes, and so on.) The anions are believed to bind to the amide groups between dendrimer chains as

18

ELECTROCHEMICALLY CONTROLLED H-BONDING

Figure 1.3 Observed DE1/2 for the ferrocene 0/ þ 1 CV wave upon addition of NBu4HSO4 to different generations of Fc dendrimers in NBu4BF4/CH2Cl2. The x-axis indicates the number of equivalents of HSO4 added per ferrocene unit.30

shown in structure 21. As the dendrimer generation grows, the chains are forced closer together creating a tighter and more preorganized binding cleft that leads to stronger binding and larger potential shifts. Similar effects have been observed with a variety of different ferrocene dendrimers, as well as Au nanoparticles that have amidoferrocene-terminated alkyl thiols attached to their surfaces.32 These type of receptors, where the binding site is a wedge in between outward radiating chains, have been classified as “exo” receptors.

O

Fc NH

Fc O NH O Fc N H O Fc NH N O

Fc H N O Fc H N Fc N N O O O N H H N O O Fc N O Fc N O H H O N O Fc N N N Fc O O O H H N O O O N Fc O N Fc H N H N N O Fc N O 20 O H Fc HN N O HN H Fc Fc O Fc

+Fc N H O O S H O O O H N +Fc O

21

19

1.6 EXAMPLES OF REDUCTION-BASED

1.6 EXAMPLES OF REDUCTION-BASED, ELECTROCHEMICALLY CONTROLLED H-BONDING SYSTEMS WITH NEUTRAL GUESTS 1.6.1

Flavins

Although the idea of using H-bonding to control redox properties is fairly recent in man’s chemistry, nature has been doing this for quite a long time. A nice example is seen in flavoenzymes that use the redox-active flavin group, 21, to catalyze a wide variety of reactions, such as oxidation of amines to imines and the hydroxylation of aromatic compounds. That one cofactor can catalyze such a range of reactions is possible because the protein environment interacts with the flavin through a variety of intermolecular interactions, including H-bonding, so as to adjust the redox chemistry to be appropriate for each particular task. R N

N

O N

N

+ e–

H

R N

O N

+ H+

R N

+ e–

N H R N

H

O

R N N H

N

O N

O

H

ð1:15Þ

O

N

N

O N

N

O 21 R N

N

H

N H

N

O N

ð1:16Þ H

O

N

O N

ð1:17Þ H

O

The reactions most commonly involved in flavin redox chemistry are shown in Equations 1.15–1.17. One-electron reduction of the flavin (Eq. 1.15) produces a relatively stable radical anion. Protonation of the radical anion produces an unstable neutral radical (Eq. 1.16), which will be rapidly reduced by another electron (Eq. 1.17) to give the flavohydroquinone anion. In an effort to better understand how the redox behavior of flavins is influenced by hydrogen bonding, Rotello and coworkers reported a study in 1995 where they looked at the electrochemistry of flavin 22 (R1 ¼ R2 ¼ Me) in the presence of diamidopyridine 23 (R ¼ Et).33 This compound hydrogen bonds to the flavin as shown, mimicking the hydrogen-bonding pattern in the active site of a number of flavoproteins. The binding constant of the diamide to the flavin in the oxidized form is 537 M1 in CDCl3 as determined by 1 H NMR. Reduction of the flavin to the radical anion would be expected to strengthen binding further by increasing the negative charge on the carbonyl oxygens, and, indeed, addition of 5 equivalents of 23 to the flavin in NBu4ClO4/CH2Cl2 resulted in a þ 155 mV shift in the E1/2 of the flavin 0/1 potential (Fig. 1.4b), indicating strong stabilization of the radical anion through H-bonding. This result was a nice confirmation of the role that

20

ELECTROCHEMICALLY CONTROLLED H-BONDING

Figure 1.4 CVs of different flavins, 22, in 0.1 M NBu4ClO4/CH2Cl2 by themselves (solid line) and in the presence of 50 mM 23 (dashed line): (a) 22, R1 ¼ Me, R2 ¼ NMe2, (b) 22, R1 ¼ R2 ¼ Me, (c) 22, R1 ¼ R2 ¼ H, and (d) 22, R1 ¼ R2 ¼ Cl. 200 mV/s scan rate.15

H-bonding plays in the flavoenzymes, but, perhaps more importantly, it really showed how strong an effect intermolecular H-bonding with a neutral binding partner could have on the redox potential of an organic redox couple, and, as a result, ushered in all the work that has followed on electrochemically controlled H-bonding with neutral guests. R O O

H N

N N

N H N

R2 R1

22

O

N 23

H N O R

21

1.6 EXAMPLES OF REDUCTION-BASED

Since their initial report on flavins, Rotello and coworkers have reported a number of other detailed studies on this system.15,34–44 One of the interesting things about the flavin/diamidopyridine system is that not only does the guest substantially alter the redox potential, but it also changes the nature of the electrode reaction. As shown in Fig. 1.4, in the absence of the guest, the flavin electrochemistry is generally not completely reversible. The return oxidation peak is too small and there is a second oxidation peak that appears at more positive potentials. Addition of the guest both shifts the reduction positive and makes it reversible. A completely reversible wave is also observed, without the guest, when the imide N is methylated, indicating that the irreversibility is due to the presence of the relatively acidic imide NH. Rotello and coworkers have explained this by suggesting that the apparent one-electron reduction of the flavin in aprotic solvent actually corresponds to the two-electron, one-proton reduction to the flavohydroquinone anion (Eqs 1.15–1.17), where the proton source is an oxidized flavin coming in from bulk solution.45 Deprotonation of the oxidized flavin prevents its reduction, with the result that the overall process is still one electron/flavin (Eq. 1.18). The return oxidation peak at more positive potentials is then due to oxidation of the flavohydroquinone anion. Addition of the diamidopyridine increases the reversibility of the flavin reduction by preventing the proton transfer both through stabilization of the imide NH on the oxidized flavin and by alteration of the charge distribution in the radical anion.15

R2

R N

2 R1

O

N N

N O

H

+

2e–

R2 R1

R N N H

O

N N O

R2

R N

H

R1

N

O

N

+

N O

21

ð1:18Þ Other groups have also looked at the electrochemistry of flavins with various Hbonding partners. One of the most interesting studies is that by Yano and coworkers who designed the elaborate binding partner 24 for the azoflavin 25.46 The pendant guanidinium groups provide four H-bonds in addition to the three set up to bond to the imide portion of the flavin. This results in very strong H-bonding even in the oxidized state, with Kox ¼ 1.9  104 M1 in 20% ACN/CH2Cl2. Reduction of the flavin to the radical anion increased the binding strength even further, giving a DE1/2¼ 317 mV. This corresponds to a binding enhancement of 2.2  105 and a Kred ¼ 4.3  109 M1, certainly one of the largest binding constants that has been reported for these types of H-bonded complexes. Because of the very strong binding in both oxidation states, this is also one of the few examples where two CV waves, one at EH and one at EHG (Scheme 1.1), are observed at less than 1 equivalent of the guest.

22

ELECTROCHEMICALLY CONTROLLED H-BONDING

H H

H

N

N H

N H

H22C12 N N

H

N

N

H N

H N

H

1.6.2

O H

25

N

N

H N

O H H

N

N N

24

N

C H N 6 13 C6H13

Arylimides

With flavins, the redox-dependent H-bonding site is the imide functional group on one side of the molecule. A year after Rotello’s initial report on the flavin/ diamidopyridine system, Smith and coworkers reported that a similar type of redox-dependent system can be created by just using a simple aromatic imide such as 1,8-naphthalimide, 26, with the same type of diamidopyridine guests.47 Subsequently, Rotello’s group has used aromatic imides in a variety of redox-dependent binding studies.48–50 Like the flavins in aprotic solvents, the aromatic imides undergo reversible reductions to radical anions in aprotic solvents, which greatly increases the strength of binding to diamidopyridine or other similarly structured binding partners. R O O N H O

H N N 23

H N O

26 R

With the flavin and imide systems, the central NH provides an additional H-bond that helps to strengthen the interaction with guest in the oxidized form. However, this is a H-donating group while the carbonyls are H-acceptors, so while reduction would be expected to increase the H-accepting ability of the carbonyls and strengthen the H-bonds between imide carbonyl and amide, it would actually be expected to decrease the strength of the H-bond between the imide NH and the pyridine N. A comparison of the E1/2 values of a flavin and its N-methyl analog in solvents of varying H-acceptor strength, along with some computational work, supports this hypothesis.51

23

1.6 EXAMPLES OF REDUCTION-BASED

Since, from the above analysis, the imide NH actually reduces the redox-dependent binding effect, an alternative way to use imides in H-bonding systems is to eliminate the NH, replacing it with an N-alkyl group.52,53 Reversible reduction still results, with now just the carbonyl oxygens being involved in the H-bonding. Since it is so straightforward to alkylate at the imide NH, this also provides a simple means to connect additional functionality. Very nice examples of this are the H-bond-based molecular shuttles reported by Leigh, Paolucci, and coworkers.54,55 Molecular shuttles are constructed from at least two separate, but physically interconnected molecular units. The assembly, called a rotaxane, consists of a long linear molecule that is “threaded” through a macrocycle. Bulky groups on either end of the “thread” prevent the macrocycle from slipping off. For a shuttle, the thread will contain at least two binding sites for the macrocycle, with one of these being strongly preferred over the other in the initial state. This preference changes upon external stimulus, such as electron transfer. O Ph Ph

N H

H N

O N

7

O

O

O

O NH

HN

NH

HN

27

O Ph Ph

N H

H N

O N

9

O

O

O

O

Ph N

28 O

O

Ph 29

The threads of the H-bonding shuttles reported by Paolucci, Leigh, and coworkers contain a succinamide station connected to either an imide, 27, or a diimide, 28, station. 1 H NMR clearly shows that the amide macrocycle, 29, strongly prefers the succinamide station in the oxidized state. However, upon reduction of the imide/ diimide, the preference changes, with now the imide/diimide station being strongly preferred. This results in net movement of the macrocycle from the succinamide to the imide/diimide (Eq. 1.19). This behavior can be deduced from the difference in the CVs of the thread by itself and that of the macrocycle-threaded shuttle. With the naphthylimide shuttle, the succinamide station is preferred over the imide 106 to 1 in the oxidized state, but, in the reduced state, the imide is preferred 500 to 1. Switching from an imide to a diimide opens up an additional control element because the diimide now undergoes two reversible reductions, first to the radical anion and then to a dianion. Analysis of the electrochemical data for this system indicates a 200 to 1 preference for the succinamide station in the zero state. This switches to a 2 to 1 preference for the diimide in the 1 state and a 4000 to 1 preference for the diimide in the 2 state.

24

1.6.3

ELECTROCHEMICALLY CONTROLLED H-BONDING

o-Quinones

In addition to the arylimides, Smith and coworkers also introduced o-quinones, discussed earlier in this chapter, as another example of a redox-dependent H-bonding receptor for neutral guests (Eq. 1.9).19,47,56 This system differs from the flavins and the imides in that a strong interaction is only observed in the reduced state, making these receptors function as on/off switches. The groups of Rotello and Cooke have also published several papers exploring the redox-dependent H-bonding properties of o-quinones.36,57–59 Many of the applications that can be envisioned for electrochemically controlled H-bonding and redox-dependent binding in general will require that at least one of the components be attached to an electrode surface. It is therefore important to determine whether behavior observed in solution can still be observed when the host is anchored to a surface. This was first tested for a redox-dependent H-bonding receptor with the o-quinone system using the phenanthrenequinone pyrrole 30.56 Electrooxidation of the pyrrole unit results in the formation of a pyrrole polymer that coats the electrode surface as it is formed. The amount of polymer deposited can be controlled by the number of CV cycles into the pyrrole oxidation wave. With 30, thick polymer layers give broad CV waves in the quinone voltage region, but thinner layers produce a well-resolved wave for the quinone 0/1 reduction, which is reasonably stable when the electrodes are placed into fresh electrolyte solution with no 30. As in solution, addition of different urea derivatives causes this wave to shift positive. The relative magnitude of the shifts mirror that seen in solution. Furthermore, the E1/2 moves back to the original potential when the derivatized electrode is put back into a blank solution containing no urea. O

O

(CH2)4 N 30

Cooke and coworkers have reported preparation of flavin-modified electrodes using a similar electropolymerization procedure.34 They have also studied electrodes coated with self-assembled monolayers (SAMs) formed from both flavin39 and phenanthrenequinone disulfides.59 The monolayers are stable in CH2Cl2 solution, and, as with the electrodes formed from 30, show redox-dependent binding behavior similar to that seen in solution. Interestingly, the phenanthrenequinone SAM

1.6 EXAMPLES OF REDUCTION-BASED

25

electrodes were also studied in the presence of phenylurea-terminated dendrimers. Addition of excess dendrimer produces a þ 200 mV shift in the E1/2 of the phenanthrenequinone 0/1 reduction, which is significantly larger than that obtained with phenanthrenequinone dissolved in solution. 1.6.4

Nitrobenzenes

Nitrobenzenes, also introduced by Smith and coworkers10 and discussed previously, provide yet another example of a simple organic redox couple in which reduction increases negative charge on two convergent oxygens, in this case the nitro oxygens (Eq. 1.3). The use of this group has not been explored as much as the flavin, imide, and o-quinone systems. However, the simplicity of its structure and the ease of which it can be introduced into organic structures suggests it may prove to be useful in the construction of more elaborate supramolecular assemblies. An example of the possible utility of the nitroaromatic group in creating more complicated structures comes from the still very simple case of 1,4-dinitrobenzene, 31.60 Almost all of the redox-dependent H-bonding systems investigated so far involve 1:1 complex formation. The ferrocene dendrimers are one exception. Another is 1,4-dinitrobenzene, where cyclic voltammetry provides very strong evidence for the formation of the 1:2 complex with diphenylurea, 2, upon reduction to the dianion (Eq. 1.19). This dianion is unusually stable because the negative charge is spread over four oxygen atoms. It is possible that it could function as a simple linker unit that could tie together other units containing, for example, two urea groups, allowing larger and more complex structures to be assembled and disassembled under redox control. Ph

Ph O

O N O

N O

+ 2e–, + 2 2

N H

O

N H

O

N

O Ph

31

1.6.5

2

O N O

H N O H N

ð1:19Þ

Ph 2

Tetrazines

All of the examples discussed above involve an increase in negative charge on two convergent oxygen atoms. Nitrogen of course is also a good H-accepting atom, but a negatively charged nitrogen would generally be expected to be quite basic, resulting in greater problems with protonation. However, recently Rotello and coworkers have reported that the tetrazine derivatives, 32 and 33, act as redox-dependent H-bonding hosts for dialkylthioureas (Eq. 1.20).611 H NMR titrations in CDCl3 indicate that there is no significant interaction with the thioureas and the tetrazines in their oxidized state. However, addition of 50 equivalents of 1-ethyl-3-octylthiourea, 34, to solutions of the tetrazines in CH2Cl2 produced a þ 80 mV shift in the E1/2 of the 0/1 reduction of 32 and a þ 60 mV shift in that of 33, indicating fairly significant interaction

26

ELECTROCHEMICALLY CONTROLLED H-BONDING

between the thioureas and the tetrazine radical anions. Computational studies support the formation of H-bonds between the urea NHs and two adjacent nitrogens in the radical anion, as indicated in Equation 1.20. X N N

X N N

+

e–

X

N N

X N N

+ 34

X

32 X = O 33 X = S

N N

Et N N

X

H N S H N C8H17

ð1:20Þ

34

1.7 EXAMPLES OF OXIDATION-BASED, ELECTROCHEMICALLY CONTROLLED H-BONDING SYSTEMS WITH NEUTRAL GUESTS 1.7.1

Ferrocene and Phenylenediamines

While using oxidation, in particular the oxidation of ferrocene to ferrocenium, has been a generally successful strategy to increase the strength of H-bonds with anionic guests, it has not as of yet been as successful with neutral H-accepting guests. Undoubtedly, part of the reason for this, as discussed earlier, is that this is such a good strategy for anions that the commonly used electrolyte anions, PF6, BF4, and ClO4, will compete for the binding site, particularly in CH2Cl2. By far, the largest DE1/2 values for oxidation-based redox-dependent H-bonding with a neutral guest reported so far are with the dimethylaminophenylurea/diamide system (Eq. 1.4).11 However, this system was studied in the presence of a large anion electrolyte. Much smaller shifts are observed with electrolytes containing PF6 and ClO4. The next largest shifts that have been reported for oxidations with neutral guests are with the ferrocene diamide, 7, reported by Tucker and coworkers. Maximum shifts of 90 mV have been reported with dicarboxylic acid guests,16 and 60 mV with cyclic ureas.62 However, these studies were done in NBu4PF6/CH2Cl2. It is possible that much larger shifts would be observed in a large anion electrolyte. 1.7.2

Tetrathiofulvalene

In addition to ferrocene, the oxidative redox couple that has received the most attention in supramolecular chemistry is tetrathiofulvalene (TTF), 35. This compound undergoes two reversible one-electron oxidations, first to a radical cation and then to a dication (Eq. 1.21). TTF first came to prominence in the 1970s when it was discovered that the charge transfer complex between it and 7,7,8,8-tetracyanoquinonedimethane (TCNQ) shows metallic conductivity. As a result, a large variety of different TTF derivatives have been prepared and characterized. This rich synthetic chemistry, coupled with the electroactivity, has intrigued supramolecular chemists for some time, with the result that the TTF unit has been incorporated into a wide variety of

27

1.8 INDIRECT ELECTROCHEMICALLY CONTROLLED H-BONDING SYSTEMS

macrocycles and other supramolecular assemblies.63 S

S

S

S

– e–

S

S

S

S

– e–

S

S

S

S

ð1:21Þ

35

Like ferrocene, TTF does not inherently contain good H-bonding functionality, but it is possible to attach good H-bonding sites that will be in conjugation with the TTF. This allows for possible perturbation of H-bond strength through an inductive-type effect. Two examples of this strategy have appeared in the literature. The first is the imide-annealed TTF derivative 36 reported by Goldenberg and Neilands,64 which like other imides will H-bond with diamidopyridines. Oxidation of the TTF should decrease the H-bonding by removing electron density from the electron-donating imide carbonyls, and indeed a modest þ 30 mV E1/2 shift for the TTF 0/ þ 1 couple is observed in NBu4PF6/CH2Cl2, indicating a threefold decrease in binding strength upon oxidation. R H O S

S

S

S 36

N O

N N

N Bu

O

H H

N

O

O R

23

S

S

S

S 37

H H N

H N 38 H N

O

Another TTF derivative modified for H-bonding that has been studied is the amidopyridine TTF, 37, prepared by Cooke, Rotello, and coworkers.65 In this case, oxidation of the TTF should increase H-bonding strength by removing electron density from the amide NH, and, in fact, a negative DE1/2 of 33 mV is observed upon addition of amide 38. Interestingly, a computational study suggests that the hydrogen adjacent to the amide group on the TTF ring also likely participates in H-bonding to the carbonyl oxygen of the guest in the þ 1 state. Note that in both of the above cases, the observed E1/2 shifts are modest, but the experiments were done in 0.1 M NBu4PF6/CH2Cl2. Much larger effects might be observed if the work was repeated using a large anion electrolyte.

1.8 INDIRECT ELECTROCHEMICALLY CONTROLLED H-BONDING SYSTEMS In the previous examples, electrochemical control of supramolecular structure is achieved by using electron transfer to directly affect H-bond strength. An alternative strategy is to not attempt to directly perturb the H-bonds, but instead to use electron

28

ELECTROCHEMICALLY CONTROLLED H-BONDING

transfer to create an additional, non-H-bonding interaction that either breaks apart or helps form a H-bonded complex. Given the large number of H-bonded complexes that have been prepared, it seems surprising that there are not more examples of this strategy. Perhaps, one of the reasons is that in this case the voltammetry is unlikely to provide unambiguous evidence that the binding is changing, and therefore alternative proof must be provided. One nice example of indirect control, for which good supporting evidence exists, is the ferrocenylurea calix[4]arene, 39, reported by Moon and Kaifer.66 It has been well established that similar urea-substituted calixarenes form dimeric capsules through H-bonds between the ureas on the two halves in noncompetitive solvents such as chloroform. 1 H NMR evidence shows that 39 also forms dimers in mixed solutions containing up to 3:10 CD3CN/CDCl3. Cyclic voltammetry of the complex in 1:10 CH3CN/CHCl3 shows a single ferrocene wave indicating that the ferrocenes are noninteracting. The thought was that oxidation of the ferrocenes to ferroceniums should break apart the capsule due to the creation of very unfavorable electrostatic interactions. However, although the CV wave is distorted from an ideal one-electron reversible process, this is not enough to prove that the capsule is breaking apart upon oxidation. More convincing evidence is that the diffusion coefficient of the oxidized calixarene, as measured by an NMR technique, is twice as fast as that of the reduced species. Furthermore, FT-IR spectroscopy shows that the peak for the urea carbonyls goes from 1657 cm1 in the reduced species, consistent with H-bonded oxygens, to 1783 and 1698 cm1, in the oxidized species, consistent with nonH-bonded oxygens. O Fc O

Fc NH O NH

NH

HN

NHHN

Fc

OR OR OR RO 39

HN

OHN

N

Fc O

N

H N

H N

O

Fe

O

H 40

H

Fe

O

N H

N H

N

O

N O

While the above example shows that electrochemically created charge can be used to break apart a H-bonded complex, this is not always the case. This is shown by a study of another ferrocene-containing H-bonded dimer, also by Kaifer and coworkers.67 The monomer in this example is based on the well-studied ureidopyrimidine framework that provides a linear AADD (A ¼ H-acceptor; D ¼ H-donor) array capable of forming four linear H-bonds with itself. The version studied by Kaifer, 40, is modified to prevent the keto-enol tautomerization that complicates the chemistry of these systems. 1 H NMR indicates that, as expected, 40 exists as a dimer in CD2Cl2, but as a monomer in the more polar CD3CN. CVs of 40 in different ratios of CH2Cl2 and CH3CN are shown in Fig. 1.5. In pure CH2Cl2, scan (a),

1.9

CONCLUSIONS

29

Figure 1.5 CVs of 1 mM 40 in 0.1 M NBu4PF6 with different ratios of CH2Cl2 and CH3CN: (a) 100% CH2Cl2, (b) 50% CH2Cl2, 50% CH3CN, (c) 20% CH2Cl2, 80% CH3CN, and (d) 100% CH3CN. 100 mV/s scan rate.67 (See the color version of this figure in Color Plates section.)

two reversible CV waves separated by 390 mV are observed. When CH3CN is added, scans (b) and (c), the first oxidation grows at the expense of the second. In pure CH3CN, scan (d), only the first wave is observed at a potential close to the first wave in CH2Cl2. These results are consistent with the NMR, with the two-wave voltammogram being due to dimer and the one-wave voltammogram being due to the monomer. The observance of two well-separated, reversible CV waves for the dimer in this system is a remarkable result. There has to be a very high degree of electronic communication between the two ferrocenes in order to make it so much harder to oxidize the second ferrocene than the first. Given the spatial distance between the ferrocenes, which are on opposite sides of the dimer, this cannot be explained by a through-space interaction, so apparently this electronic interaction is going through the H-bonds. Another surprising feature is that the second wave is reversible. This means that the two monomers dimerize in CH2Cl2 even when both are positively charged. Evidently, for this system under these experimental conditions, the four Hbonds are strong enough to overcome the electrostatic repulsion between the monomers.

1.9

CONCLUSIONS

H-bonds, being one of the strongest intermolecular interactions, are used extensively in the supramolecular chemistry of both man and nature. Their high electrostatic character makes it straightforward to perturb their strength electrochemically. There are two main ways to do this. Reduction can be used to make a H-acceptor a better acceptor by increasing the negative charge on a H-accepting atom, or oxidation can be used to make a H-donor a better donor by increasing the positive

30

ELECTROCHEMICALLY CONTROLLED H-BONDING

charge on a H-donating atom. There are now a number of examples demonstrating the effectiveness of both strategies. In a noncompetitive solvent such as CH2Cl2, it appears possible to change binding strengths by factors of 102–104, and even 105 if additional interactions such as ion–ion can be invoked. The minimum elements required to achieve these large effects appear to be that there are two strong, almost linear H-bonds between host and guest and that at least one of these is strongly perturbed by the electron transfer reaction. Based on the existing examples, the largest effects are observed if the affected H-bonds are an integral part of the redox couple, but significant effects are also possible when the redox couple is simply attached through conjugation with the H-bonding site. Other factors to take into consideration are the solvent and the electrolyte. For reductions, electrolytes with NBu4 þ as the cation appear to be good choices. For oxidations, especially with neutral guests, the best choices may be electrolytes with very large anions such as B(C6F5)4, since the more common electrolyte anions such as PF6, ClO4, and BF4 may cause interference. Another factor to be aware of is the possibility of proton transfer between host and guest. Looking toward the future, some basic work in electrochemically controlled Hbonding still remains to be done. In particular, the role of electrolyte needs to be investigated further. The number of different redox couples that have been used successfully is also still rather limited. Nonetheless, it is also clear that much of the groundwork has been laid, and that the next major steps will be to couple what has been learned about how electrochemistry can be used to control the strength of Hbonds in simpler systems with advances in H-bond assembly in solution, in order to construct larger, more sophisticated supramolecular structures operating under electrochemical control. Leigh and Paolucci’s work with molecular shuttles and Kaifer’s work with redox-active capsules are steps in this direction, but this is really just the beginning.

REFERENCES L. J. Prins, D. N. Reinhoudt, P. Timmerman, Angew. Chem., Int. Ed. 2001, 40, 2382. R. P. Sijbesma, E. W. Meijer, Chem. Commun. 2003, 5. P. S. Corbin, S. C. Zimmerman, J. Am. Chem. Soc. 2000, 122, 3779. K. A. Jolliffe, P. Timmerman, D. N. Reinhoudt, Angew. Chem., Int. Ed. 1999, 38, 933. J. RebekJr., Angew. Chem., Int. Ed. 2005, 44, 2068. F. Hof, S. L. Craig, C. Nuckolls, J. RebekJr., Angew. Chem., Int. Ed. 2002, 41, 1488. D. T. Bong, T. D. Clark, J. R. Granja, M. R. Ghadiri, Angew. Chem., Int. Ed. 2001, 40, 988. C. A. Nijhuis, B. J. Ravoo, J. Huskens, D. N. Reinhoudt, Coord. Chem. Rev. 2007, 251, 1761. 9. G. A. Jeffrey, An Introduction to Hydrogen Bonding, Oxford University Press, New York, 1997. 10. J. Bu, N. D. Lilienthal, J. E. Woods, C. E. Nohrden, K. T. Hoang, D. Truong, D. K. Smith, J. Am. Chem. Soc. 2005, 127, 6423. 1. 2. 3. 4. 5. 6. 7. 8.

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11. J. E. Woods, Y. Ge, D. K. Smith, J. Am. Chem. Soc. 2008, 130, 10070. 12. S. R. Miller, D. A. Gustowski, Z. H. Chen, G. W. Gokel, L. Echegoyen, A. E. Kaifer, Anal. Chem. 1988, 60, 2021. 13. M. Gomez, F. J. Gonzalez, I. Gonzalez, Electroanalysis 2003, 15, 635. 14. M. Gomez, F. J. Gonzalez, I. Gonzalez, J. Electrochem. Soc. 2003, 150, E527. 15. Y.-M. Legrand, M. Gray, G. Cooke, V. M. Rotello, J. Am. Chem. Soc. 2003, 125, 15789. 16. J. D. Carr, S. J. Coles, M. B. Hursthouse, M. E. Light, J. H. R. Tucker, J. Westwood, Angew. Chem., Int. Ed. 2000, 39, 3296. 17. N. Gupta, H. Linschitz, J. Am. Chem. Soc. 1997, 119, 6384. 18. J. E. Woods, D. K. Smith, unpublished results. 19. Y. Ge, L. Miller, T. Ouimet, D. K. Smith, J. Org. Chem. 2000, 65, 8831. 20. R. J. LeSuer, W. E. Geiger, Angew. Chem., Int. Ed. 2000, 39, 248. 21. F. Barriere, W. E. Geiger, J. Am. Chem. Soc. 2006, 128, 3980. 22. P. D. Beer, Acc. Chem. Res. 1998, 31, 71. 23. P. D. Beer, P. A. Gale, Angew. Chem., Int. Ed. 2001, 40, 486. 24. C. H. Park, H. E. Simmons, J. Am. Chem. Soc. 1968, 90, 2431. 25. P. D. Beer, A. D. Keefe, J. Organomet. Chem. 1989, 375, C40. 26. P. D. Beer, D. Hesek, J. Hodacova, S. E. Stokes, J. Chem. Soc., Chem. Commun. 1992, 270. 27. P. D. Beer, C. Hazlewood, D. Hesek, J. Hodacova, S. E. Stokes, J. Chem. Soc., Dalton Trans. 1993, 1327. 28. P. D. Beer, M. G. B. Drew, A. R. Graydon, D. K. Smith, S. E. Stokes, J. Chem. Soc., Dalton Trans. 1995, 403. 29. F. Oton, A. Tarraga, A. Espinosa, M. D. Velasco, P. Molina, J. Org. Chem. 2006, 71, 4590. 30. C. Valerio, J.-L. Fillaut, J. Ruiz, J. Guittard, J.-C. Blais, D. Astruc, J. Am. Chem. Soc. 1997, 119, 2588. 31. D. Astruc, M.-C. Daniel, J. Ruiz, Chem. Commun. 2004, 2637. 32. M.-C. Daniel, J. Ruiz, S. Nlate, J.-C. Blais, D. Astruc, J. Am. Chem. Soc. 2003, 125, 2617. 33. E. Breinlinger, A. Niemz, V. M. Rotello, J. Am. Chem. Soc. 1995, 117, 5379. 34. G. Cooke, J. Garety, S. Mabruk, V. Rotello, G. Surpateanu, P. Woisel, Chem. Commun. 2004, 2722. 35. J. B. Carroll, G. Cooke, J. F. Garety, B. J. Jordan, S. Mabruk, V. M. Rotello, Chem. Commun. 2005, 3838. 36. G. Cooke, H. A. de Cremiers, F. M. A. Duclairoir, J. Leonardi, G. Rosair, V. M. Rotello, Tetrahedron 2003, 59, 3341. 37. E. C. Breinlinger, V. M. Rotello, J. Am. Chem. Soc. 1997, 119, 1165. 38. A. K. Boal, V. M. Rotello, J. Am. Chem. Soc. 1999, 121, 4914. 39. G. Cooke, F. M. A. Duclairoir, P. John, N. Polwart, V. M. Rotello, Chem. Commun. 2003, 2468. 40. G. Cooke, J. F. Garety, B. Jordan, N. Kryvokhyzha, A. Parkin, G. Rabani, V. M. Rotello, Org. Lett. 2006, 8, 2297. 41. C. Bourgel, A. S. F. Boyd, G. Cooke, H. Augier de Cremiers, F. M. A. Duclairoir, V. M. Rotello, Chem. Commun. 2001, 1954. 42. G. Cooke, Y.-M. Legrand, M. Rotello Vincent, Chem. Commun. (Camb.) 2004, 1088.

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ELECTROCHEMICALLY CONTROLLED H-BONDING

43. M. Gray, A. J. Goodman, J. B. Carroll, K. Bardon, M. Markey, G. Cooke, V. M. Rotello, Org. Lett. 2004, 6, 385. 44. A. S. F. Boyd, J. B. Carroll, G. Cooke, J. F. Garety, B. J. Jordan, S. Mabruk, G. Rosair, V. M. Rotello, Chem. Commun. 2005, 2468. 45. A. Niemz, J. Imbriglio, V. M. Rotello, J. Am. Chem. Soc. 1997, 119, 887. 46. T. Kajiki, H. Moriya, K. Hoshino, T. Kuroi, S. I. Kondo, T. Nabeshima, Y. Yano, J. Org. Chem. 1999, 64, 9679. 47. Y. Ge, R. R. Lilienthal, D. K. Smith, J. Am. Chem. Soc. 1996, 118, 3976. 48. R. Deans, A. Niemz, E. C. Breinlinger, V. M. Rotello, J. Am. Chem. Soc. 1997, 119, 10863. 49. A. Niemz, V. M. Rotello, J. Am. Chem. Soc. 1997, 119, 6833. 50. M. Gray, A. O. Cuello, G. Cooke, V. M. Rotello, J. Am. Chem. Soc. 2003, 125, 7882. 51. A. O. Cuello, C. M. McIntosh, V. M. Rotello, J. Am. Chem. Soc. 2000, 122, 3517. 52. J. B. Carroll, M. Gray, K. A. McMenimen, D. G. Hamilton, V. M. Rotello, Org. Lett. 2003, 5, 3177. 53. S. I. Kato, T. Matsumoto, K. Ideta, T. Shimasaki, K. Goto, T. Shinmyozu, J. Org. Chem. 2006, 71, 4723. 54. A. Altieri, F. G. Gatti, E. R. Kay, D. A. Leigh, D. Martel, F. Paolucci, A. M. Z. Slawin, J. K. Y. Wong, J. Am. Chem. Soc. 2003, 125, 8644. 55. G. Fioravanti, N. Haraszkiewicz, E. R. Kay, S. M. Mendoza, C. Bruno, M. Marcaccio, P. G. Wiering, F. Paolucci, P. Rudolf, A. M. Brouwer, D. A. Leigh, J. Am. Chem. Soc. 2008, 130, 2593. 56. Y. Ge, D. K. Smith, Anal. Chem. 2000, 72, 1860. 57. G. Cooke, V. Sindelar, V. M. Rotello, Chem. Commun. 2003, 752. 58. J. B. Carroll, M. Gray, G. Cooke, V. M. Rotello, Chem. Commun. 2004, 442. 59. G. Cooke, J. Couet, J. F. Garety, C.-Q. Ma, S. Mabruk, G. Rabani, V. M. Rotello, V. Sindelar, P. Woisel, Tetrahedron Lett. 2006, 47, 3763. 60. C. Chan-Leonor, S. L. Martin, D. K. Smith, J. Org. Chem. 2005, 70, 10817. 61. B. J. Jordan, M. A. Pollier, L. A. Miller, C. Tiernan, G. Clavier, P. Audebert, V. M. Rotello, Org. Lett. 2007, 9, 2835. 62. J. Westwood, S. J. Coles, S. R. Collinson, G. Gasser, S. J. Green, M. B. Hursthouse, M. E. Light, J. H. R. Tucker, Organometallics 2004, 23, 946. 63. M. B. Nielsen, C. Lomholt, J. Becher, Chem. Soc. Rev. 2000, 29, 153. 64. L. M. Goldenberg, O. Neilands, J. Electroanal. Chem. 1999, 463, 212. 65. A. S. F. Boyd, G. Cooke, F. M. A. Duclairoir, V. M. Rotello, Tetrahedron Lett. 2003, 44, 303. 66. K. Moon, A. E. Kaifer, J. Am. Chem. Soc. 2004, 126, 15016. 67. H. Sun, J. Steeb, A. E. Kaifer, J. Am. Chem. Soc. 2006, 128, 2820.

CHAPTER 2

Molecular Motions Driven by Transition Metal Redox Couples: Ion Translocation and Assembling–Disassembling of Dinuclear Double-Strand Helicates VALERIA AMENDOLA and LUIGI FABBRIZZI Dipartimento di Chimica Generale, Universita di Pavia, Pavia, Italy

2.1

INTRODUCTION

There exists a current interest in the design of molecular systems capable of converting chemical energy into an intramolecular movement, thus producing mechanical work. These systems have been defined as molecular machines and their synthesis and investigation represent one of the most active and appealing areas of supramolecular chemistry.1 The energy necessary for promoting the motion is provided by an ancillary reaction: when this reaction has a redox nature, being carried out either chemically or electrochemically, we have the molecular equivalent of the electrical machines of the macroscopic world. In order to perform controlled and reversible movements and to behave as a machine, the envisaged molecular system should have a mobile and a fixed component: one of the components should be redox active and the oxidized and reduced states should have almost comparable stability and should be connected by a reversible, and possibly fast, electron transfer process. The two oxidation states should display a different topological affinity with respect to the other component, so that a redox change can induce a modification of the topology of the whole molecular system, generating an intramolecular motion. The occurrence of fast and reversible movements also requires that the interaction between the mobile and the fixed part is based on Electrochemistry of Functional Supramolecular Systems. Edited by Paola Ceroni, Alberto Credi, and Margherita Venturi Copyright  2010 John Wiley & Sons, Inc.

33

34

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

a labile interaction: electrostatic, hydrogen bonding, p–p donor–acceptor, and metal– ligand (involving labile metal centers). First redox-driven intramolecular movements (and first electrochemical machines) were reported in 1994 and involved a rotaxane held by p–p interactions2 and a catenane based on metal–ligand interactions.3 In the Stoddart’s rotaxane (1),2 the wheel (A, conventionally assumed as the mobile component) had a strong p-acceptor nature, while the axle (the fixed component) contained two unequivalent stations of varying p-donor properties: a 4,40 -benzidine (D1, more p-donor) and 4,4-dimethoxydiphenyl (D2, less p-donor). Under unperturbed conditions, A stays on D1. However, D1 is redox active, undergoing fully reversible one-electron oxidation to D1 þ . Thus, on chemical or electrochemical oxidation to D1 þ , the station loses its donor tendencies and the wheel A conveniently moves to D2. Then, on addition of a proper reducing agent or adjusting the working electrode to the appropriate cathodic potential, D1 þ reduces to D1 and A moves back to the more donating station, resetting the system. In this machine, the D1/D1 þ couple is the engine and the auxiliary redox reaction provides the fuel.

Another redox-driven intramolecular movement involved the half-turn of one ring of an asymmetric catenane.3 The Sauvage’s catenane consisted of two intertwined rings, one containing a phen fragment (2) and the other containing both a phen and a terpy fragment (3).

The two interlocked rings are held together through the coordinative interactions to a copper center. When copper is in the þ 1 oxidation state, it is coordinated by two phen subunits, one from each ring. Such a coordinative arrangement is determined by the preference of the posttransition cation CuI (d10 electronic configuration) toward four-coordination, according to a tetrahedral geometry. On one-electron oxidation, the transition cation CuII (electronic configuration d9) forms, which tends to increase its coordination number, to 5 in the present case, through the coordination of the phen subunit of ring 2 and the terpy subunit of ring 3. This promotes a half-turn of ring 3, thus

2.1

INTRODUCTION

35

Figure 2.1 A square scheme illustrating a redox-driven intramolecular motion. Species with an asterisk (Ox* and Red*) are metastable and tend to rearrange to their stable topological isomer (Ox and Red).

giving rise to a controlled intramolecular motion. On CuII-to-CuI reduction, the halfturn is reversed and the system resets. In the present case, the engine of the electrical machine is represented by the CuI/CuII couple, fueled by an auxiliary redox reduction. The pioneering papers by Stoddart and Sauvage have stimulated the design of a variety of movable rotaxanes and catenanes, whose controlled motion is promoted by a redox change. In all cases, the process of the redox-driven intramolecular motion can be described by a square scheme, as illustrated in Fig. 2.1. It is assumed that the species Red at the top left corner of the scheme is the stable form of the supramolecular system, under unperturbed conditions (e.g., rotaxane 1, with the wheel placed on the benzidine station).2 On one-electron oxidation (step (i)), a metastable species is generated, Ox* (e.g., 1 þ , with the wheel still located on the benzidinium radical cation fragment), which thermally evolves to the stable form Ox, according to the “translational” step (ii). On reduction of Ox (step (iii)), the metastable species Red* forms (1 with the wheel on the dimethoxydiphenyl station), which undergoes thermal rearrangement to Red, according to a reverse “translational” step. The direct movement consists of the consecutive steps (i) and (ii) and its rate is determined by the slowest step. Electron transfer processes are in general fast; thus, the rate of the intramolecular motion is in most cases controlled by the “translational” step (ii). This step is thermal in nature and is affected by a kinetic barrier, whose magnitude is related to the steric interactions between the mobile and the fixed component. Sliding of wheel A from D1 to D2 and the D2-to-D1 reverse motion are fast processes, probably due to the flexibility of the aliphatic segment connecting the two stations. On the other hand, in the Sauvage’s catenate,3 the translational step (ii), which follows the CuI-to-CuII oxidation, can take hours. This may be due to the fact that the half-turn of ring 3 primarily involves the dissociation of one or more CuIIN bonds, a rather endothermic process due to the loss of ligand field (LF) energy experienced by the transition cation CuII. On the other hand, the reverse movement is fast because the d10 metal center CuI does not lose LF energy on CuIN bond dissociation. The square scheme illustrated in Fig. 2.1 applies to any redox process followed by a conformational rearrangement and was first discussed in detail considering the redox-driven linkage isomerization of the sulfoxide ligand associated with the RuIIRuIII change in [RuII,III(NH3)5(sulfoxide)]2 þ /3 þ complexes (S-coordination to RuII and O-coordination to RuIII).4 There exist other types of redox-driven intramolecular motions that can be interpreted on the basis of the square scheme of Fig. 2.1 and are promoted by

36

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

a metal-centered redox couple. They include, among the others, (a) the translocation of ions (either metal ions or anions) within ditopic ligands and receptors and (b) the assembling and disassembling of double helices. Representative examples of these processes will be described in the following sections. 2.2 2.2.1

ION TRANSLOCATION Metal Translocation

In the most general situation, a redox-active metal ion is translocated from a given site to another site of the same molecular system, following a chemical (a redox reaction) or an electrochemical input. The redox-driven reversible translocation of a metal ion in a two-component molecular system is schematically sketched in Fig. 2.2. The box in Fig. 2.2 symbolizes a system containing two distinct and separate ligating compartments, A and B, which differ in their coordinating properties. In particular, A is a hard receptor and B a soft one. Given a metal M possessing two adjacent oxidation states of comparable stability, Mn þ of soft nature and M(n þ 1) þ hard in character, in a solution containing equimolecular amounts of the reduced metal Mn þ and the two-compartment system A–B, Mn þ will occupy the soft compartment B. But, if Mn þ is oxidized, chemically or electrochemically, the hard ion that forms, M(n þ 1) þ , will find it thermodynamically convenient to move from B to the nearby hard compartment A. Consecutive oxidation and reduction processes would make the metal center M shuttle back and forth, between A and B, along a defined route. The square scheme of Fig. 2.1 applies to the translocation process. In particular, the rates of the translocation processes should depend (i) on the nature of the coordinative bonds between M and receptors A and B, whether kinetically labile or inert, and (ii) on the feasibility of the conformational rearrangement associated with the metal displacement. In the present case, the metal center plays two roles: that of the mobile part and that of the engine, fueled by the reaction with an oxidizing or a reducing agent added to the solution. Two redox couples that fit well the prerequisites mentioned earlier are FeIII/FeII and CuII/CuI. 2.2.1.1 Metal Translocation Based on the FeIII/FeII Couple The first example of redox-driven translocation of a metal center was based on the FeIII/FeII couple and took place in ditopic ligands containing (i) a tris-hydroxamate compartment and (ii) a tris-(2,20 -bipyridine) compartment.5

Figure 2.2

The redox switched translocation of a metal ion within a two-compartment ligand.

2.2

37

ION TRANSLOCATION

The bidentate hydroxamate (4), containing two oxygen donor atoms and having a negative charge, can be considered a hard ligand and forms very stable 1 : 3 octahedral complexes with tripositive metal ions such as FeIII (d5 electronic configuration, highspin state). The log b3 value for the [FeIIIL] (L ¼ 4, R1 ¼ CH3, R2 ¼ H, acetohydroxamate) in water is 28.33.6 It is suggested that metal coordination stabilizes the mesomeric form 4b, in which both oxygen donor atoms retain a negative charge. On the other hand, the bpy ligand gives rise to especially stable [MII(bpy)3]2 þ complexes of octahedral geometry with the FeII cation (d6, low spin). The high stability is due to the extended back-donation from the filled dp orbitals of the metal to the low-energy p* molecular orbital of bpy molecules. The two-compartment ligand 5 consists of three strands, each containing a hydroxamate subunit and a bpy subunit, covalently linked to a tripodal platform. In the system illustrated in Fig. 2.3, the platform is provided by the tetramine tren.

N N

N

N N

N

N

N N

N

O

NH O

O

HN

O

O N

O

O

O O

+ e–

O

O

NH

HN

NH

O

O

–e – N

HO

N

HN N OH

HO

O NH

O

HN

N

5a (FeIII)

O

N N

O NH

O

O

NH

NH

HN

O

N

5b (FeIII)

Figure 2.3 Translocation of an iron center within a two-compartment ligand, driven by the FeIII/FeII redox couple. FeIII prefers the inner compartment, which provides six oxygen donor atoms and retains a triply negative charge; FeII chooses the peripheral compartment consisting of three bpy subunits. Consecutive chemical reduction and oxidation makes the metal move back and forth between the two compartments.

38

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

On addition of 1 equivalent of FeIII to a MeOH/H2O solution of the ditopic ligand 5, the solution takes a light brown color with formation of a rather intense band centered at 420 nm (e ¼ 2400 M1 cm1). These optical features pertain to FeIII tris-hydroxamate complexes and show that the metal ion has occupied the tris-hydroxamate compartment. On addition of ascorbic acid, the solution turns violet, while an intense band centered at 540 nm develops (e ¼ 4700 M1 cm1). Such a band results from a metal-to-ligand charge transfer transition, typically occurring in the [FeII(bpy)3]2 þ chromophore. Thus, the FeIII-to-FeII reduction process promotes the translocation of the iron center from the inner compartment (complex 5a) to the peripheral one (5b), as pictorially illustrated in Fig. 2.3. The 5a-to-5b translocation process has a lifetime t ¼ 22 s. The reverse process can be induced on oxidation of the FeII center with S2O82, but it is remarkably slower, taking place over a period of minutes, at 70 C! The sluggishness of the process must be probably ascribed to the complexity of the conformational changes experienced by the three strands. On the other hand, the extreme slowness of the 5a-to-5b translocation may be due to the intrinsic sluggishness of the redox processes involving peroxydisulfate ion as an oxidizing agent.

Iron translocation based on the FeIII/FeII redox couple has been later observed in the ditopic ligand 6,7 which is structurally similar to 5. In analogy with the ditopic ligand 5, system 6 consists of three equivalent strands, covalently linked to a tren platform. Each strand contains a salicylamide group and a bpy subunit. Salicylamide easily deprotonates to give a bidentate ligand whose oxygen donor atoms retain a negative charge, delocalized on both through a p mechanism (see mesomeric forms 7a and 7b).

2.2

ION TRANSLOCATION

39

Compartment A consists of three salicylamide groups: each group behaves as a negatively charged ligand and offers two oxygen donor atoms. Thus, compartment A exhibits a definitely hard nature and is expected to display a specific affinity toward high-spin FeIII. Compartment B, like in the previously considered system 5, consists of three 2,20 -bipyridine fragments and establishes strong coordinative interactions with low-spin FeII. The 1:1 complex of FeIII and 6, in a DMF/H2O solution, shows an orange color (lmax ¼ 460 nm, e ¼ 3200 M1 cm1), as expected for a tris-salicylamide FeIII complex, thus indicating that FeIII resides in compartment A. On addition of a reducing agent (e.g., ascorbic acid), the solution gradually turns violet, while an absorption band develops (lmax ¼ 574 nm, e ¼ 2300 M1 cm1, with a shoulder at 543, 2250). These spectral features are those expected for a [FeII(bpy)3]2 þ chromophore and demonstrate that, on reduction, the metal center has moved to compartment B. On addition, at room temperature, of the strong oxidizing agent H2O2, a violet solution of the FeII complex again takes an orange color, due to the formation of the FeIII tris-salicylaldimine complex. Hydrogen peroxide provides on reduction hydroxide ions (H2O2 þ 2e ¼ 2OH), which favor the deprotonation of the OH group of each salicylate moiety and consequent coordination of the phenolate oxygen atom to FeIII. Kinetic studies showed that translocation processes in either direction take place on the 10 s timescale. The higher rate of the A-to-B translocation process can be ascribed to the greater flexibility of system 6 with respect to 5. Moreover, the use of the very reactive oxidizing agent H2O2, compared to S2O82, may account for the much faster back-translocation occurring within the ditopic receptor 6 with respect to 5. A FeIII/FeII-driven translocation process of different nature is illustrated in Fig. 2.4.8 The ligand 8 consists of a 4-methylphenol platform, to which two different terdentate subunits have been appended in 2- and 6-positions. One appendance consists of a tertiary amine nitrogen atom and two phenolate oxygen atoms (deprotonation of all the phenolic groups of 8 is guaranteed by the presence of the base 2,4,6trimethylpyridine (collidine) in the MeCN solution). The other appendance possesses one tertiary amine nitrogen atom and two pyridine nitrogen atoms. When 1 equivalent of FeIII(ClO4)3 is added to a MeCN solution of 8, in the presence of collidine, the FeIII cation seeks for the coordination of the three available oxygen atoms, one from the platform and two from the NO2 appendance. Six-coordination, to give the preferred octahedral geometrical arrangement, probably distorted because of the steric constraints, is achieved by coordination of the nitrogen atom of the NO2 appendance and two MeCN molecules (indicated with S in formula 8a in Fig. 2.4). On the other hand, if stoichiometric FeII(ClO4)2 is added to a solution of 8 containing collidine, the FeII center looks for the coordination of the three nitrogen atoms of the N3 appendance, while the phenolate oxygen atom of the platform and two solvent molecules complete

40

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

Figure 2.4 The “pendular” motion of an iron center. The left compartment provides a donor set consisting of three negatively charged phenolate oxygen atoms, a tertiary amine nitrogen atom, and two solvent molecules (S ¼ MeCN) and hosts FeIII. The right compartment offers as donors a phenolate oxygen atom, one tertiary amine nitrogen atom, two pyridine nitrogen atoms, and two solvent molecules and is suitable for the coordination of FeII. The central phenolate oxygen atom is shared by the two compartments and acts as a pivot in the translocation process.

six-coordination. The above evidence provides the basis for the occurrence of a redox-driven translocation of the iron center between the two ligating compartments, which share the central phenolate oxygen atom. Indeed, translocation was carried out electrochemically and was investigated through cyclic voltammetry (CV) experiments at a platinum working electrode, in a MeCN solution made 0.1 M in [Bu4N]ClO4. The CV experiment (Fig. 2.5) starts at a potential of 800 mV (a), where the FeII ion stays in the right compartment (8b in Fig. 2.4), which provides the donor set [N3O þ 2S]. On increasing the potential, oxidation to FeIII takes place, with

b FeIII

Current

oxidation

–0.8

c

–0.4

0

Potential (V)

a

reduction

FeII d

Figure 2.5 Cyclic voltammogram of a MeCN solution equimolar (4  103 M) in 8, FeII(ClO4)2, and collidine. Working electrode: platinum disk (5 mm diameter); supporting electrolyte: [Bu4N]ClO4; scan rate: 0.1 V/s; reference electrode: AgNO3/Ag in MeCN. Diagram adapted from Ref. 8.

2.2

ION TRANSLOCATION

41

Figure 2.6 A square scheme illustrating the “pendular” motion of an iron center, driven by the FeII/FeIII redox couple. As judged from voltammetric experiment carried out at varying potential scan rate, the lifetime t for both translocation processes is CuII > NiII and can be monitored through voltammetric titration experiments. In particular, a differential pulse voltammetry (DPV) scan was carried out on a MeCN solution 103 M in the [CuII(11)NiII](ClO4)4 complex salt and 0.1 M in [Bu4N]ClO4, using a microsphere platinum electrode. The [CuII–NiII]4 þ system II

Figure 2.9 Redox-driven translocation of the anion X (e.g., chloride), based on the NiII/NiIII change. The nickel center acts both as an engine and as a receptor for the X anion (when in the NiIII state). Occurrence of the reversible X translocation is afforded by the following sequence of anion affinity: NiIII > CuII > NiII.

2.2

ION TRANSLOCATION

45

Figure 2.10 Differential pulse voltammetry profiles obtained at a platinum working electrode in a MeCN solution 5  104 in [Et3Bn]Cl and 0.1 M in [Bu4N]ClO4. Dashed line, 103 M in the [CuII(11)NiII](ClO4)4 complex salt; solid line, 103 M in both the [CuII(12)](ClO4)2 and [NiII(11)](ClO4)2 complex salts.

undergoes a one-electron oxidation at 0.74 V versus Fc þ /Fc, which corresponds to the NiII-to-NiIII change. On addition of Cl, a new peak develops at a much less positive potential (0.24 V), while the intensity of the peak at 0.74 V progressively decreases. In particular, the peak intensity at 0.24 V reaches a plateau with the addition of 1 equivalent of Cl; further anion addition does not induce the appearance of any other peak. Figure 2.10 (dashed line) displays the DPV profile recorded on addition of 0.5 equivalent of chloride. Such a voltammetric behavior can be explained on assuming that on NiII-to-NiIII oxidation a Cl ion leaves the CuII center and occupies an axial position on the NiIII center. In particular, the voltammetric response is based on the thermodynamic cycle illustrated in Fig. 2.11.

(2.1)

(2.2) Figure 2.11 Thermodynamic cycle that connects equilibrium (2.1) (the NiII-to-NiIII redox change and chloride ion translocation) and equilibrium (2.2) (simple NiII-to-NiIII oxidation) through chloride complexation equilibria involving CuII and NiIII, in the conjugate system [CuII(11)NiII]4 þ .

46

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

Note that the horizontal equilibrium (2.1) in the cycle corresponds to the translocation process pictorially illustrated in Fig. 2.9. Moreover, F  ½E ðClÞE  ¼ F  DEðVÞ ¼ RT  ½log KðNiIII --ClÞ--log KðCuII --ClÞ As DE ¼ 0.5 V and log K(CuIICl) ¼ 5.7, at 25 C, log K(NiIIICl) ¼ 8.5, which fulfills the first requirement for anion translocation. On the other hand, 23.08  DE(V) gives the DG value (in kcal/mol) for the “neat” translocation equilibrium (2.3): ð2:3Þ In equilibrium (2.3), the chloride ion is transferred from the CuII center to the already oxidized NiIII cation. Such an equilibrium is characterized by a very favorable free energy change DG ¼ 11.5 kcal/mol. At this stage, one could argue whether process (2.3) is authentically intramolecular (i.e., the chloride which has moved on NiIII is really coming from CuII) or intermolecular (the chloride leaves CuII for the solution, while a different chloride ion from the solution binds NiIII).

In this connection, it would be useful to consider a similar process that is unquestionably intermolecular. In particular, we considered the complex of CuII with “component” ligand 12 (a N-benzyl-substituted tren derivative) and the NiII complex with the other “component” ligand 13 (a N-benzyl-substituted cyclam derivative). The N-benzyl substituent is there in order to reproduce in the separated complexes the coordinating environments present in the conjugate system 11. Then, we carried out a voltammetric titration on a MeCN solution 103 M both in [CuII(12)]2 þ and in [NiII(13)]2 þ . Prior to chloride addition, an oxidation peak appeared at 0.74 V, which corresponds to the NiII/NiIII couple within macrocycle 13 and superimposes exactly with the peak observed for the [CuII(11)NiII]4 þ complex. On chloride addition, the peak intensity at 0.74 V decreases, while a new peak develops at 0.57 V. Figure 2.10 displays the DPV profile obtained for a solution containing 0.5 equivalent of [Et3Bn]Cl (solid line). It has to be noted that the NiII-to-NiIII oxidation takes place at a much more positive potential than in the case of the [CuII(11)NiII]4 þ complex. ð2:4Þ

2.2

ION TRANSLOCATION

47

From peak separation DE (0.17 V), it is possible to calculate the DG value corresponding to the intermolecular equilibrium (2.4), in which a chloride ion is detached from the [CuII(12)]2 þ complex and goes through the solution to coordinate the metal center in the [NiIII(13)]3 þ complex (DG ¼ 3.9 kcal/mol), a quantity significantly less negative than observed for the “neat” translocation equilibrium (2.3). In particular, the “intramolecular” anion translocation equilibrium (2.3) is favored by an extra free energy of 7.6 kcal/mol with respect to the intermolecular anion transfer between the separated components. An interesting question is whether such an energy advantage has an enthalpic or an entropic origin. Very conveniently, DS values associated with the “neat” anion translocation equilibria (2.3) and (2.4) can be determined by carrying out voltammetric investigations at varying temperature. In fact, DE ¼ 

DH  DS þ T F F

ð2:5Þ

Thus, DE should vary linearly with the temperature and DS could be calculated from the slope of the linear plot DE versus T. DPV experiments on the conjugate system [CuII(11)NiII]4 þ ( þ 0.5 equivalent of Cl) and on system {[CuII(12)]2 þ [NiII(13)]2 þ } ( þ 0.5 equivalent of Cl) were carried out over the 20 to 30 C temperature range. Pertinent DE versus T plots are shown in Fig. 2.12. It is observed that DE(3), for the conjugate complex [CuII(11)NiII]4 þ (open triangles in Fig. 2.12), does not vary along the investigated temperature interval, indicating that DS ¼ 0. Thus, DG ¼ DH ¼ 11.5 kcal/mol, which indicates that the

Figure 2.12 Temperature dependence of DE for MeCN solutions of (i) [CuII(11)NiII]4 þ (open triangles) and (ii) an equimolar mixture of [CuII(12)]2 þ and [NiII(13)]2 þ (filled triangles), in the presence of 0.5 equivalent of Cl. DE is the peak separation in the DPV profiles in Fig. 2.10 and refers to the neat anion translocation equilibria (2.3) and (2.4). The slope of the DE versus temperature straight line gives DS /F for pertinent equilibrium.

48

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

favorable free energy change associated with the neat translocation equilibrium (2.6) is solely due to a very exothermic thermal contribution. This reflects the especially high energy of the NiIIICl interaction within the cyclam subunit compared to the CuIICl bond in the tren moiety of the conjugate system [CuII(11)NiII]4 þ . On the other hand, the DE value measured for the {[CuII(12)]2 þ þ [NiII(13)]2 þ } system shows a significant linear decrease with the increasing temperature (filled triangles in Fig. 2.12), which indicates a negative value of the entropy change. In particular, from the slope of the least squares straight line the following value was calculated: DS ¼ 28 cal/mol/K. The algebraic combination of DG (3.9 kcal/mol) and TDS (8.3 kcal/mol, at 25 C) gives DH ¼ 12.2 kcal/mol, a value comparable to that pertaining to the neat anion translocation equilibrium involving the conjugate system (DH ¼ 11.5 kcal/mol). The enthalpy changes typically reflect the balance of bonding energy terms: the closeness of DH values for equilibria (2.3) and (2.4) suggests that the balance of the bonding terms (breaking of the CuIICl bond and forming of the NiIIICl bond) is nearly the same for both the conjugate system [CuII(11)NiII]4 þ and the two separate components [CuII(12)]2 þ and [NiII(13)]2 þ . It derives that the significant disadvantage experienced by the neat anion translocation process between separate components solely results from a negative entropy effect. It is suggested that such an effect has a probabilistic nature. In this connection, one should consider that anion translocation results from the collision of the CuII(tren)(Cl) þ subunit with the NiIII center both in the conjugate and in the separate component system. It is intuitive that the process is more probable in the covalently linked system [CuII(11)NiII]4 þ than in the {[CuII(12)]2 þ þ [NiII(13)]2 þ } mixture. In fact, in the [CuII(11)NiII]4 þ complex, useful collisions are mainly intramolecular and result from the occasional folding of the covalently linked system, which brings CuII(tren)(Cl) þ and NiIII(cyclam)3 þ subunits into contact. On the other hand, for separate components, collisions involve complexes dispersed in the solution and the probability of their occurrence is related to the concentration. The greater the concentration, the higher the probability that the two complexes collide. A semiquantitative comparison can be attempted in the following way. The distance between copper and nickel centers in the [CuII(11)NiII]4 þ complex  is ca. 7.5 A, as calculated through molecular modeling. Thus, we can assume that within a sphere whose center is represented by NiIII and the CuIICl fragment moves  3 whose volume is 1766 A , that is, 1.766  1024 L. Thus, the concentration of the active CuIICl fragment is 1.766  1024 molecule/L, or, in the more familiar molar scale, 0.94 M. This concentration is 103–104 higher than that of the [CuII(12)]2 þ complex in the solution electrochemically or spectrophotometrically investigated and accounts for the higher probability (and the lower loss of entropy) of the intramolecular chloride translocation in the [CuII(11)NiIII]5 þ system with respect to intermolecular anion transfer from [CuII(13)Cl] þ to [NiIII(13)]3 þ . There are not many other anions undergoing a CuII-to-NiIII translocation within the conjugate system 11. Reducing anions such as Br, I, and NCS undergo oxidation at a potential lower or comparable to that of the NiII/NiIII couple, which prevents any translocation. NO3 and HSO4 are resistant to oxidation, but show a moderate affinity toward the CuII center. In particular, in a solution containing equimolar

2.3

ASSEMBLING/DISASSEMBLING OF HELICATE COMPLEXES

49

amounts of the dimetallic system [CuII(11)NiII]4 þ and the anion, only a small fraction of the anions is bound to the CuII center. Among inorganic anions relevant to coordination chemistry, only NCO displays a “regular” behavior. In fact, it is quite resistant to the oxidation and gives a stable complex with the reduced form of the covalently linked system, [CuII(NCO)(11)NiII]3 þ , with log K ¼ 4.4 0.1. It derives that, in a solution 103 M both in [CuII(11)NiIII]4 þ and in NCO, 82% of the anions are bound to CuII. Then, over the course of a DPV titration experiment, a new peak develops at 0.27 V versus Fc þ /Fc and reaches a plateau after the addition of 1 equivalent of NCO (while the peak at 0.74 V decreases and disappears). Also in the present case the DE value (0.47 V) is remarkably larger than DE measured for a solution containing the separate components (0.26 V). This substantiates the observation that the intramolecular anion translocation from CuII to NiIII within the covalently linked system [CuII(11)NiIII]5 þ is distinctly favored with respect to the intermolecular translocation between the separate components. The engine of the investigated “machine” is provided by the NiII/NiIII redox change. One could ask whether the ancillary metal center (CuII) could be replaced by another one. In this connection, we considered that ZnII tetramine complexes exhibit a definite tendency to form five-coordinate complexes with a given anion X. In particular, they show a pronounced affinity toward the carboxylate donor group. Spectrophotometric titration experiments with the benzoate anion, BzO, in a solution of the [ZnII(11)NiII]4 þ system in MeCN/MeOH (1:1 v/v) indicated the formation of a rather stable [ZnII(BzO)(11)NiII]3 þ complex (log K ¼ 5.6 0.1). However, DPV studies showed that benzoate addition neither altered the intensity of the peak at 0.74 V versus Fc þ /Fc (pertaining to NiII-to-NiIII oxidation) nor induced a shift of the peak. This suggests that the benzoate anion forms a more stable complex with the ZnII(tren)2 þ subunit than with the NiIII(cyclam)3 þ fragment. Thus, the required sequence of anion binding tendencies (NiIII > MII > NiII) is not fulfilled and the redox-induced anion translocation process cannot occur.

2.3 ASSEMBLING/DISASSEMBLING OF HELICATE COMPLEXES DRIVEN BY THE CuI/CuII COUPLE Molecule 14 is a tetradentate ligand consisting of two bpy moieties linked in 3- and 30 -positions by a CH2OCH2 spacer; the plain bpy ligand forms complexes of formula [MI(bpy)2] þ of tetrahedral geometry with d10 metal ions, such as CuI and AgI, whose extra stability results from the donation of electron density from filled dp orbitals of the metal to empty p* molecular orbitals of bpy molecules. One would expect that mixing of 1 equivalent of 14 with 1 equivalent of [CuI(MeCN)4]ClO4 in a MeCN solution would give the 1:1 complex, in which 14 acts as a quadridentate ligand and fully chelates the d10 metal center. However, this does not occur due to the steric constraints, which prevents the achievement of the tetrahedral coordination geometry required by the CuI center. Instead, metal and ligand find it more convenient to give a complex of 2 : 2 stoichiometry, [Cu2I(14)2]2 þ ,

50

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

Figure 2.13 The dinucleating bis-bidentate ligand 14 forms with MI metal ions of electronic configuration d10 (e.g., CuI, AgI) dimetallic complexes of formula [M2I(14)2]2 þ , in which two molecules of 14 are intertwined to give a double helix. Ligands of the type 14 are named helicands and complexes such as 15 are called helicates. In this particular case, we have a double-strand helicate.

in which each CuI center is coordinated by one half of each bis-bidentate ligand and can achieve the preferred tetrahedral coordination geometry without not too serious strain.16 The structure of the dicopper(I) complex is outlined in Fig. 2.13 (15): the two molecules of 14 are arranged as a double helix. Following the language of coordination chemistry, molecule 14 has been named a helicand and complex 15 a helicate. Since the seminal papers by Lehn,17 the chemistry of helicates has greatly developed to give systems of higher complexity, involving a larger number of metals of different coordination numbers and geometry. However, the aspect of helicate chemistry relevant to this chapter refers to the redox behavior of dinuclear copper(I) double-strand helicate complexes, in particular to the consequences of the CuI-to-CuII oxidation process on a [Cu2I(L \ L)2]2 þ complex, where L \ L represents a bisbidentate ligand such as 14. This aspect has been electrochemically investigated by several authors.18 In this chapter, we will consider a homogeneous series of systems studied in our laboratory, in order to offer a complete view of the topic.19 To begin, the dinucleating helicand 16 will be considered.

16 gives stable complexes with CuI and CuII, both in MeCN solution and in the solid state. In particular, the two complex salts [Cu2I(16)2](CF3SO3)2, dinuclear, and [CuII(16)](CF3SO3)2, mononuclear, were isolated in the crystalline form and corresponding crystal and molecular structures were determined through X-ray diffraction

2.3

ASSEMBLING/DISASSEMBLING OF HELICATE COMPLEXES

51

Figure 2.14 The molecular structure of the [Cu2I(16)2]2 þ double-strand helicate complex cation. CuI metal centers are represented as spheres. The hydrogen atoms of the two strands have been omitted for clarity. Structure redrawn from data deposited at the Cambridge Crystallographic Data Centre: CCDC 118958.

studies. Figure 2.14 shows the double-strand helix arrangement of the bis-bidentate ligand 16 in the [Cu2I(16)2]2 þ complex cation. Each CuI center (gray ball) is bound to an imine and to a pyridine nitrogen atom from each strand and shows a rather distorted tetrahedral coordination geometry. On the other hand, Fig. 2.15, which displays the molecular structure of the [CuII(16)](CF3SO3)2 salt, shows that the CuII ion prefers to form a mononuclear complex species. In particular, the transition metal ion CuII finds it convenient to reach tetragonal coordination through the chelation by a single molecule of 16, in order to better benefit from ligand field stabilization energy terms. On assuming that the geometrical features described above are maintained in solution, the CuII/CuI redox change in solution would result in an assembling–disassembling equilibrium, as pictorially illustrated in Fig. 2.16. The occurrence of the redox-driven reversible assembling–disassembling process involving copper complexes of 16 has been verified through cyclic voltammetry experiments at a platinum electrode in a MeCN solution. Figure 2.17 shows the CV profile obtained with a solution of the double-strand helicate complex [Cu2I(16)2]2 þ . The starting potential has been set at 300 mV (versus Fc þ /Fc), where the dinuclear complex is stable (a). On increasing the potential, a poorly defined oxidation peak develops (b), corresponding to the CuI-to-CuII oxidation process. At 900 mV, the potential is reversed. However, going down to 500 mV and less, the reduction peak opposite to peak b is not observed. In fact, the oxidation to CuII is followed by a fast

52

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

Figure 2.15 The molecular structure of the [CuII(16)]2 þ . The hydrogen atoms of ligand 16 have been omitted. The CuII center experiences a rather distorted square coordination geometry. Structure redrawn from data deposited at the Cambridge Crystallographic Data Centre: CCDC 118957.

disassembling process to give two mononuclear [CuII(16)]2 þ complexes. The CuII mononuclear complex, due to the preferred tetragonal coordination geometry, is especially stable to the reduction, which takes place at a much less positive potential (peak d, at 0 mV). Following reduction, the two CuI mononuclear complexes quickly reassemble to give the double-strand helicate species, closing the cycle. The corresponding square scheme is outlined in Fig. 2.18. The high irreversibility of the profile depends upon the fact that the assembling– disassembling process is too fast with respect to the timescale of the CV experiment

Figure 2.16 The redox-driven disassembling of a dicopper(I) double-strand helicate complex to give two mononuclear copper(II) complexes, in which each strand behaves as a quadridentate ligand. On subsequent reduction, the two mononuclear complexes reassemble to give the helicate. The illustrated process fits well the behavior of copper complexes of 16 in a MeCN solution.

2.3

ASSEMBLING/DISASSEMBLING OF HELICATE COMPLEXES

53

Figure 2.17 Cyclic voltammogram of a MeCN solution of [CuII(16)](CF3SO3)2. Supporting electrolyte: 0.1 M [Bu4N]ClO4; scan rate: 0.1 V/s; internal reference electrode: Fc þ /Fc. Diagram adapted from Ref. 20.

(the profile in Fig. 2.17 has been taken at a potential scan rate of 100 mV/s). However, the same irreversible profile was observed also at a potential scan rate of 20 V/s. Considering that CV experiments of the type illustrated in Fig. 2.17 cover an interval of ca. 1 V, it derives that the transient species [Cu2II(16)]4 þ (upper right corner in the square scheme of Fig. 2.18) and [CuI(16)] þ (lower left corner of square scheme) have a lifetime lower than 50 ms.

Figure 2.18 A square scheme illustrating the disassembling of the [Cu2I(16)2]2 þ double helicate complex, following CuI-to-CuII oxidation, and the consequent assembling of two [CuII(16)]2 þ mononuclear complexes, following the CuII-to-CuI reduction. The process ultimately derives from the geometrical coordinative preferences of the two oxidation states: CuI prefers a tetrahedral coordination, which can be achieved with the double helicate arrangement; CuII prefers a square coordination geometry, which is fulfilled by the coordination of a single molecule of 16.

54

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

However, it has been recently demonstrated by Pallavicini et al. that the lifetime of the dicopper(II) double-strand helicate [Cu2II(16)]4 þ can be significantly increased by introducing hindering substituents on the framework of 16. In particular, this was shown to occur with the copper complexes of the bis-bidentate ligand 17.21

The bidentate ligand 17 forms bis with CuI a double-strand helicate complex, whose structure was elucidated through X-ray diffraction studies and is shown in Fig. 2.19. The [Cu2I(17)2]2 þ complex displays geometrical features very similar to those of the [Cu2I(16)2]2 þ analogue: as an example, the CuICuI distance is 3.71 A, to be compared to 3.67 A. The CV behavior of the [Cu2I(17)2]2 þ complex was investigated in a MeCN solution made 0.1 M in [Bu4N]ClO4, using a platinum microsphere as a working

Figure 2.19 The molecular structure of the [Cu2I(17)2]2 þ double-strand helicate complex cation. CuI metal centers are represented as spheres. Hydrogen atoms of the two strands have been omitted for clarity. Structure redrawn from data deposited at the Cambridge Crystallographic Data Centre: CCDC 641164.

2.3

ASSEMBLING/DISASSEMBLING OF HELICATE COMPLEXES

55

Figure 2.20 Cyclic voltammogram of a MeCN solution of [Cu2I(17)2]2 þ double-strand helicate complex. Supporting electrolyte: [Bu4N]ClO4; scan rate: 0.2 V/s; internal reference electrode: Fc þ /Fc. Diagram adapted from Ref. 21.

electrode. Figure 2.20 shows the pertinent profile: the potential was set at 0.00 V versus the reference couple (Fc þ /Fc) and was scanned toward positive values. Very interestingly, two consecutive waves were observed, with peaks located at 190 and 420 mV. The first wave originates from the one-electron oxidation of one CuI center (to CuII) and the second wave from the one-electron oxidation of the second CuI ion. Peak separation DE ¼ 230 mV originates from the repulsive electrostatic effect exerted by the CuII ion on the oxidation of the proximate CuI center and from the statistical effect (36 mV). On the reverse scan, two distinct and well-shaped reduction waves develop at 330 and 90 mV, respectively (peak separation 240 mV). This unambiguously indicates that on oxidation of the dicopper(I) helicate a dicopper(II) helicate forms according to two distinct one-electron oxidation steps. The [Cu2II(17)2]4 þ helicate complex is stable over the timescale of the CV experiment, even when performed at a potential scan rate as low as 20 mV/s. This indicates for the dicopper(II) helicate a lifetime of 50 s or more. Thus, the voltammetric behavior can be simply interpreted on the basis of the following stepwise equilibria, of mere electrochemical nature, involving intact double-strand helicates: ½Cu2 I ð17Þ2 2 þ K ½CuI CuII ð17Þ2 3 þ þ e

ð2:6Þ

½CuI CuII ð17Þ2 3 þ K ½CuII CuII ð17Þ2 4 þ þ e

ð2:7Þ

On the other hand, the CV investigation on a MeCN solution containing the mononuclear [CuII(17)]2 þ complex gave totally different results. Pertinent profile is shown in Fig. 2.21.

56

MOLECULAR MOTIONS DRIVEN BY TRANSITION METAL REDOX COUPLES

Figure 2.21 Cyclic voltammogram of a MeCN solution of [CuII(17)]2 þ mononuclear complex. Supporting electrolyte: 0.1 M [Bu4N]ClO4; scan rate: 0.2 V/s; internal reference electrode: Fc þ /Fc. Diagram adapted from Ref. 21.

The starting potential was again set at 850 mV and the potential was then scanned toward negative values. In particular, on moving toward cathodic potentials, a rather intense peak was observed at 150 mV, which, on the reverse scan, was not restored, indicating irreversible behavior. On the other hand, on moving further toward more positive potentials, two consecutive well-defined waves were observed to develop, with peaks at 160 and 380 mV, respectively (peak separation 220 mV). The voltammetric behavior can be explained as follows: (i) the reduction of the mononuclear [CuII(17)]2 þ complex to the corresponding mononuclear [CuI(17)] þ species is immediately followed by a fast assembling to give the dicopper(I) double-strand helicate. On scanning the potential to positive values, the [Cu2I(17)2]2 þ complex undergoes stepwise one-electron oxidation processes to give first [CuICuII(17)2]3 þ and then [CuIICuII(17)2]4 þ . The appearance of the two well-shaped one-electron peaks indicates that the double helix arrangement is maintained in the CuII oxidation state. The above evidence indicates that the OCH3 substituent in each quinoline moiety slows down the disassembling of the CuII double-strand helicate complex, which persists in solution on the timescale of the CV experiment (in the present case performed at a scan rate of 200 mV/s). This kinetic effect has probably to be related to the fact that methoxy substituents raise the energy of the transition state, which must involve a planarization of the tetrahedral coordinative arrangement within the CuII helicate complex. Any thermodynamic effect has to be ruled out, otherwise the mononuclear [CuII(17)]2 þ complex will tend to assemble spontaneously to the helicate complex [Cu2II(17)2]4 þ , even prior to the reduction. Double-strand dicopper helicate complexes are interesting systems in that they may show hysteresis (as observed with ligand 16), thus giving rise to a rare example

REFERENCES

57

of electrochemical bistability. It has also been shown that the hysteretic behavior can be controlled and modulated through simple synthetic modifications, as observed with system 17.

2.4

CONCLUDING REMARKS

In this chapter, we have considered different types of molecular movements, which are driven by metal-centered redox couples: FeIII/Fe/II and CuII/CuI. Due to the rich redox activity of transition metals, one would expect that the choice could be extended to other redox couples, within the d block of the periodic table. However, the simple occurrence of a fast and reversible one-electron redox process of a transition metal does not suffice for promoting molecular motions or conformational changes in the appropriate molecular system. In fact, it is required that the two oxidation states of the metal center exhibit distinctly different electronic and/or geometrical preferences. This is evident for the CuII-to-CuI change, in which the border is crossed between the realms of transition and posttransition metals, an almost unique opportunity (unless one would explore the d0/d1 change, perhaps TiIV/TiIII). The FeIII-to-FeII reduction process works well because it implies a change from a high-spin d5 cation, which cannot take advantage from LFSE effects and is therefore inclined to establish mainly electrostatic interactions, to a low-spin d6 cation, which benefits most from LFSE in an octahedral coordinative environment and, in the presence of p-acceptors ligands, can also exert back-donation. A d5-to-d6 change is observed in the RuIII/RuII couple, but in most complexes the second row RuIII center is present in the low-spin state, a condition that reduces its difference with respect to RuII. Thus, it seems probable that the design of electrochemically fueled engines based on metal centers will keep relying on the two couples described in the previous examples. This may further stimulate, rather than depress, the ingenuity of synthetic chemists.

REFERENCES 1. (a) V. Balzani, A. Credi, F. M. Raymo, J. F. Stoddart, Angew. Chem. 2000, 112, 3484–3530; Angew. Chem., Int. Ed. 2000, 39, 3348–3391. (b) Special issue on Molecular Machines (Ed.: J. F. Stoddart), Acc. Chem. Res. 2001, 34. (c) Special volume on Molecular Machines and Motors (Ed.: J.-P. Sauvage), Struct. Bond. 2001, 99. (d) A. H. Flood, R. J. A. Ramirez, W. Q. Deng, R. P. Muller, W. A. Goddard, J. F. Stoddart, Aust. J. Chem. 2004, 57, 301–322. (e) Special volume on Molecular Machines (Ed.: T. R. Kelly), Top. Curr. Chem. 2005, 262. (f) J.-P. Sauvage, Chem. Commun. 2005, 1507–1510. (g) K. Kinbara, T. Aida, Chem. Rev. 2005, 105, 1377–1400. (h) G. S. Kottas, L. I. Clarke, D. Horinek, J. Michl, Chem. Rev. 2005, 105, 1281–1376. (i) W. R. Browne, B. L. Feringa, Nat. Nanotechnol. 2006, 1, 25–35. (j) E. R. Kay, D. A. Leigh, F. Zerbetto, Angew. Chem. 2007, 119, 72–196;Angew. Chem., Int. Ed.2007, 46, 72–191;(k) Special issue on Molecular Machines and Switches (Eds.: A. Credi, H. Tian), Adv. Funct. Mater. 2007, 17. (l) V. Balzani, A. Credi, M. Venturi, ChemPhysChem 2008, 9, 202–220. (m) V. Balzani, A. Credi, M. Venturi, Molecular Devices and Machines, 2nd edition, Wiley-VCH, Weinheim, 2008.

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2. R. A. Bissell, E. Co´rdova, A. E. Kaifer, J. F. Stoddart, Nature 1994, 369, 133–137. 3. A. Livoreil, C. O. Dietrich-Buchecker, J.-P. Sauvage, J. Am. Chem. Soc. 1994, 116, 9399–9400. 4. A. Tomita, M. Sano, Inorg. Chem. 1994, 33, 5825–5830. 5. L. Zelikovich, J. Libman, A. Shanzer, Nature 1995, 374, 790–792. 6. R. J. Motekaitis, Y. Sun, A. E. Martell, Inorg. Chem. 1991, 30, 1554–1556. 7. T. R. Ward, A. Lutz, S. P. Parel, J. Ensling, P. G€utlich, P. Buglyo´, C. Orvig, Inorg. Chem. 1999, 38, 5007–5017. 8. P. Belle, J.-L. Pierre, E. Saint-Aman, New J. Chem. 1998, 1399–1402. 9. J.-M. Lehn, Supramolecular Chemistry. Concepts and Perspectives, VCH, Weinheim, 1995, pp. 134–135. 10. V. Amendola, L. Fabbrizzi, C. Mangano, P. Pallavicini, Acc. Chem. Res. 2001, 34, 488–493. 11. C. Canevet, J. Libman, A. Shanzer, Angew. Chem., Int. Ed. 1996, 35, 2657–2670. 12. S. Zahn, J. W. Canary, Angew. Chem., Int. Ed. 1998, 37, 305–307. 13. G. De Santis, L. Fabbrizzi, D. Iacopino, P. Pallavicini, A. Perotti, A. Poggi, Inorg. Chem. 1997, 36, 827–832. 14. L. Fabbrizzi, F. Gatti, P. Pallavicini, E. Zambarbieri, Chem. Eur. J. 1999, 5, 682–690. 15. L. Sabatini, L. Fabbrizzi, Inorg. Chem. 1979, 18, 438–444. 16. J.-M. Lehn, A. Rigault, J. Siegel, J. Harrowfield, B. Chevrier, D. Moras, Proc. Natl. Acad. Sci. USA 1997, 84, 2565–2569. 17. J.-M. Lehn, Supramolecular Chemistry. Concepts and Perspectives, VCH, Weinheim, 1995, p. 146. 18. (a) J.-P. Gisselbrecht, M. Gross, J.-M. Lehn, J.-P. Sauvage, R. Ziessel, C. PiccinniLeopardi, J. M. Arrieta, G. Germain, M. Van Meerssche, Nouv. J. Chim. 1984, 8, 661–667. (b) Y. Yao, M. W. Perkovic, D. P. Rillema, C. Woods, Inorg. Chem. 1992, 31, 3956–3962. (c) K. T. Potts, K. M. Keshavarz, F. S. Tham, H. D. Abrun˜a, C. R. Arana, Inorg. Chem. 1993, 32, 4422–4435. (d) K. T. Potts, K. M. Keshavarz, F. S. Tham, H. D. Abrun˜a, C. Arana, Inorg. Chem. 1993, 32, 4436–4449. (e) K. T. Potts, K. M. Keshavarz, F. S. Tham, H. D. Abrun˜a, C. R. Arana, Inorg. Chem. 1993, 32, 4450–4456. (f) R. Ziessel, A. Harriman, J. Suffert, M. T. Youinou, A. De Cian, J. Fischer, Angew. Chem., Int. Ed. Engl. 1997, 36, 2509–2511. (g) M. Greenwald, M. Eassa, E. Katz, I. Willner, Y. Cohen, J. Electroanal. Chem. 1997, 434, 77–82. 19. (a) V. Amendola, L. Fabbrizzi, C. Mangano, P. Pallavicini, E. Roboli, M. Zema, Inorg. Chem. 2000, 39, 5803–5806. (b) V. Amendola, L. Fabbrizzi, P. Pallavicini, Coord. Chem. Rev. 2001, 435, 216–217. (c) V. Amendola, L. Fabbrizzi, L. Gianelli, C. Maggi, C. Mangano, P. Pallavicini, M. Zema, Inorg. Chem. 2001, 40, 3579–3587. (d) V. Amendola, L. Fabbrizzi, P. Pallavicini, E. Sartirana, A. Taglietti, Inorg. Chem. 2003, 42, 1632–1636. (e) V. Amendola, L. Fabbrizzi, E. Mundum, P. Pallavicini, Dalton Trans. 2003, 773–774. 20. V. Amendola, L. Fabbrizzi, L. Linati, C. Mangano, P. Pallavicini, V. Pedrazzini, M. Zema, Chem. Eur. J. 1999, 5, 3679–3688. 21. P. Pallavicini, M. Boiocchi, G. Dacarro, C. Mangano, New J. Chem. 2007, 31, 927–935.

CHAPTER 3

Molecular Encapsulation of Redox-active Guests ANGEL E. KAIFER Center for Supramolecular Science, Department of Chemistry, University of Miami, Coral Gables, FL, USA

3.1

INTRODUCTION

The dictionary definition of the verb encapsulate is “to place in or as if in a capsule.” In chemistry, molecular encapsulation is widely understood as the placing of a molecule inside a larger one. This definition immediately suggests the use of supramolecular interactions to achieve the encapsulated state, that is, we can take advantage of intermolecular forces to create some sort of self-assembled complex in which the smaller molecule (usually referred to as the guest) is held inside the cavity of the larger molecule (the host). If the encapsulated guest is held inside the host during an experimentally meaningful period of time, we can say that the host encapsulates the guest. While this is indeed widely accepted terminology, molecular encapsulation is also commonly used in reference to other situations, which may be quite different. For instance, it is entirely possible to covalently attach one or more branching polymeric structures to a functional group, resulting in a macromolecule (usually a dendrimer) containing the functional group at the core, partially or fully surrounded by dendritic material.1 The result is the site isolation of the functional group, that is, its relative isolation or protection from solvent molecules or other solutes in the medium. However, this situation is often also referred to as molecular encapsulation of the functional group. This general case, better described as covalent encapsulation, is treated in detail in a different chapter. Therefore, we will not consider it here in any significant detail. Recent advances in materials science have made possible the preparation of microscopic particles (in nanometer or micrometer size scales) containing molecules Electrochemistry of Functional Supramolecular Systems. Edited by Paola Ceroni, Alberto Credi, and Margherita Venturi Copyright  2010 John Wiley & Sons, Inc.

59

60

MOLECULAR ENCAPSULATION OF REDOX-ACTIVE GUESTS

coated with a different material, such as silica. However, this review will focus on well-characterized molecular materials and we will leave aside the study of these nano- or microparticle based materials, which are not discrete molecular systems. Thus, the primary focus of this chapter revolves around the idea of noncovalent encapsulation,2–7 in which a small molecule, the guest, is spontaneously included inside a larger host molecule that contains a preformed or guest-induced cavity. Guest inclusion by the host gives rise to a stable inclusion complex, which represents the encapsulated state. Of course, given the tenor of this volume, we will focus on redox-active guests. More specifically, we will pay particular attention to moieties capable of fast electron transfer reactions, since these are particularly amenable for electrochemical measurements.

3.2

THERMODYNAMIC AND KINETIC CONSIDERATIONS

At the microscopic level, the formation of a host–guest inclusion complex is a dynamic process. Usually, a combination of intermolecular forces will bring together the host and the guest to form the complex. These interactions may include ion–ion, ion–dipole, hydrogen boding, van de Waals, p–p stacking, and solvophobic forces, among others, and their combined effect affords the complex a degree of stability. In thermodynamic terms, the equilibrium between the host (H) and the guest (G) to form the complex (HG) is H þ G L HG

ð3:1Þ

The free energy change (DGo) associated with this process is a measure of the relative stability of the HG complex. For a complex that forms spontaneously, that is, self-assembles in the solution phase, DGo < 0. The free energy is related to the association equilibrium constant K by the well-known thermodynamic equation: DGo ¼ RT ln K

ð3:2Þ

Furthermore, the equilibrium constant is also related to the equilibrium activities of the species involved. For convenience, the activities are often replaced by concentrations to yield the more practical expression K¼

½HGeq ½Heq ½Geq

ð3:3Þ

In biochemistry, it is much more common to write the equilibrium as a complex dissociation process and use instead the dissociation equilibrium constant (Kd), whose value is simply the reciprocal of the association K value. The biochemists’ preference to handle Kd values has a pragmatic advantage, related to the fact that Kd has concentration units. Consider, for instance, a complex association equilibrium for

3.2

THERMODYNAMIC AND KINETIC CONSIDERATIONS

61

which K ¼ 106 M1. The same process can be represented by a dissociation equilibrium constant of Kd ¼ 106 M, which represents the lower end of host (and guest) concentration levels required for the equilibrium to be effective in the association direction. In other words, if Kd ¼ 106 M (or K ¼ 106 M1), at least micromolar concentrations of host and guest must be mixed for a significant fraction of complex to be formed. This equilibrium will not be effective at forming the HG complex when submicromolar concentrations of H and G are used. Therefore, it is common to speak of millimolar, micromolar, or nanomolar equilibrium constants depending on their magnitude. An association equilibrium constant K (also referred to as a binding constant) with a value of 5  103 M1 will not be effective at guest and host concentrations of 10 mM. In fact, simple equilibrium calculations reveal that ca. 95% of the host and the guest will remain unassociated under those conditions. These straightforward arguments are unfortunately some times forgotten when dealing with inclusion complex formation. In addition to this, the dynamic character of the equilibrium is also important. If we could follow the behavior of individual guest molecules randomly moving in solution, we would see that, every once in a while, they collide with a host molecule. Some times, these collisions are productive and lead to the formation of an inclusion complex, in which the guest is encapsulated by the host (Scheme 3.1). However, the inclusion complex is stable for a limited period of time and, eventually, undergoes dissociation to regenerate the free host and guest. The equilibrium association constant K can also be expressed as the ratio of the two kinetic rate constants shown in Scheme 3.1, that is, K¼

kIN kOUT

ð3:4Þ

The kinetic rate constant for the association process (kIN) has an upper limit set by diffusion. In other words, the rate of the fastest association processes cannot exceed the rate by which the host and the guest diffuse to encounter in solution. The maximum value of kD can then be estimated using the well-known Smoluchowski equation8: kD ¼ 1000  ½4prHG ðDH þ DH ÞNA 

ð3:5Þ

Scheme 3.1 Definition of the two kinetic rate constants relevant to the formation of a host– guest inclusion complex.

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MOLECULAR ENCAPSULATION OF REDOX-ACTIVE GUESTS

where rHG is the approach distance at which H and G react, DH and DG are the corresponding diffusion coefficients, and NA is Avogadro’s number. In aqueous solution, the upper limit for kD is in the range 109–1010 M1 s1. However, most host– guest complexation reactions are significantly slower since there are many relative molecular orientations in which host–guest collisions do not lend themselves to complex formation. Also, the host’s cavity entrance (portal) may be relatively constricted, in such a way that even with a suitable relative orientation, a collision may not be productive. Therefore, it is quite common to observe kIN values several orders of magnitude under the diffusion limit given by Equation 3.5. If the kIN value and the binding constant K are known, we can use Equation 3.4 to calculate the value of kOUT; the lifetime of the complex is simply the reciprocal of kOUT. Let us take as an example a host–guest complexation process with a typical kIN value of ca. 106 M1 s1. If the binding constant K is around 103 M1, the kOUT value should be around 103 s1, which translates to a complex lifetime of ca. 1 ms. These values can be considered representative of inclusion complexation by cyclodextrin hosts. On the other hand, if the binding constant reaches a larger value of 106 M1, the complex lifetime approaches a longer value of 1 s, assuming that kIN remains approximately the same. This would be the typical situation with a cucurbituril inclusion complex. In fact, as the stability of the inclusion complex increases, the steric fit and degree of complementarity between host and guest normally improve, which may lead to lower kIN values. Overall, higher thermodynamic stability often goes hand in hand with longer complex lifetimes. This may have important consequences on the mechanistic details of electron transfer reactions, as will be addressed later in this chapter.

3.3 CYCLODEXTRIN AND CUCURBITURIL COMPLEXATION OF REDOX-ACTIVE GUESTS One of the simplest forms of guest encapsulation is the formation of an inclusion complex with a host containing a well-defined, preformed cavity. In this regard, the cyclodextrins9–11 and the cucurbit[n]urils12–14 are probably the most relevant watersoluble hosts, and both are readily accessible with a variety of cavity sizes. While the extent of encapsulation that can be achieved with these host families is usually incomplete, their wide accessibility and aqueous solubility has led to their frequent use to complex redox-active guests. Therefore, it seems appropriate to provide here a brief description of both host families, emphasizing their relative advantages and disadvantages. Both cyclodextrins (CDs) and cucurbit[n]urils (CBns) are relatively water-soluble molecules that contain a rather rigid, well-defined cavity whose inner surface is best described as hydrophobic. Both host families are capable of forming inclusion complexes in aqueous media with a variety of guests. Although the guests are not limited to those containing hydrophobic moieties, the most stable inclusion complexes are typically formed by guests with a hydrophobic residue that fits well inside the host cavity. In spite of these general similarities between CDs and CBns, it is also important to establish clearly some of their most important differences. The CDs

3.3

CYCLODEXTRIN AND CUCURBITURIL COMPLEXATION OF REDOX-ACTIVE GUESTS

63

TABLE 3.1 Structures of Unmodified Cucurbit[n]uril (Left) and Cyclodextrin (Right) Receptors and Relevant Parameters for the Most Representative Members of These Host Families

Host

n

Molecular Weight

Cavity  Diameter (A)

Aqueous Solution (mM)

CB6 CB7 CB8 a-CD b-CD g-CD

6 7 8 6 7 8

996 1163 1329 972 1135 1297

3.9–5.8 5.4–7.3 6.9–8.8 4.7–5.3 6.0–6.5 7.5–8.3

0.02 20