Non-Covalent Interactions: Theory and Experiment: Volume 2 (Theoretical and Computational Chemistry Series) - Hardcover

Pavel Hobza

Hardcover

ISBN 10: 1847558534 ISBN 13: 9781847558534

Publisher: Royal Society of Chemistry, 2009

Synopsis

Co-authored by an experimentalist (Klaus M³ller-Dethlefs ) and theoretician (Pavel Hobza), the aim of this book is to provide a general introduction into the science behind non-covalent interactions and molecular complexes using some important experimental and theoretical methods and approaches.

About this title

About the Author

Professor Dr Pavel Hobza is at the Institute of Organic Chemistry and Biochemistry, Academy of Sciences of the Czech Republic and Center for Biomolecules and Complex Molecular Systems, Prague, Czech Republic.

Professor Klaus Müller-Dethlefs is Chair of Chemistry and Chair of Physics, and Director of the Photon Science Institute, The University of Manchester, Manchester, UK.

From the Back Cover

The aim of this book is to provide a general introduction into the science behind non-covalent interactions and molecular complexes using some important experimental and theoretical methods and approaches. It is the first monograph on this subject written in close collaboration between a theoretician and an experimentalist which presents a coherent description of non-covalent interactions viewed from these two perspectives. The book describes the experimental and theoretical techniques, and some results obtained by these, which are useful in conveying the principles underlying the observable or computable properties of molecular clusters. The chemical and physical background underlying non-covalent interactions are treated comprehensively and non-covalent interactions is contrasted to ionic, covalent and metallic bonding. The role of dispersion and electrostatic interactions, static and induced multipole moments, charge transfer and charge localisation and de-localisation are described. In addition, the nomenclature and classification of non-covalent interactions and molecular clusters is discussed since there is still no unique agreement on it. The authors were among first who coined the term non-covalent for intermolecular interactions and all interactions can thus be categorised as metallic, covalent and non-covalent. The book covers covalent bonding where the properties of a moiety in a molecular cluster are concerned, for instance its electrostatic multipole moments. The historic development of the field is also briefly outlined, starting from van der Waals who first recognized the fact that molecules in the gas phase interact, through London who explained the fact that non-polar uncharged systems attract each other, making a connection to modern work of theoreticians and experimentalists who have contributed to the present knowledge in the field. The role of non-covalent interactions in nature is discussed and the book also argues why non-covalent interactions and not covalent ones play a key role in biological systems. The authors show the unique significance of non-covalent interactions in biological systems and describe several important processes (molecular recognition, structure of biomacromolecules, etc) that are fundamentally determined by non-covalent interactions. The book is aimed at undergraduate and graduate students who need to learn more about non-covalent interactions and their role in chemistry, physics and biology. It also provides valuable information to non-specialist scientists and also those who work in the area who will find it interesting reading. As both experimental and theoretical procedures are covered, this enables the reader to orientate themselves in this very intensely growing area

Excerpt. © Reprinted by permission. All rights reserved.

Non-covalent Interactions

Theory and Experiment

By Pavel Hobza, Klaus Müller-Dethlefs

The Royal Society of Chemistry

Copyright © 2010 Pavel Hobza and Klaus Müller-Dethlefs
All rights reserved.

ISBN: 978-1-84755-853-4

Contents

Chapter 1 Introduction, 1, 
Chapter 2 Characteristics of Non-covalent Complexes and Their Determination by Experimental and Theoretical Techniques, 21, 
Chapter 3 Potential-Energy and Free-Energy Surfaces, 70, 
Chapter 4 Classification of Non-covalent Complexes, 134, 
Chapter 5 Interpretation of Experimental Results and Types of Molecular Clusters, 155, 
Chapter 6 Extended Molecular Clusters in Chemistry, the Atmosphere and Stereospecific Molecular Recognition, 206, 
Subject Index, 221, 

CHAPTER 1

Introduction

1.1 An Historical Remark

It was none other than van der Waals, in the 1870s, who realised that the discrepancies observed between the state function of a real gas and the ideal gas law could be accounted for by the attracting forces between molecules or rare gas atoms. Van der Waals introduced an equation of state suitable for describing the behaviour of real (in contrast to ideal) gases. Although this law does not provide the most accurate functional description for a real gas, it nevertheless constituted a major breakthrough. Van der Waals made it explicitly obvious, for instance with respect to condensation of all real gases, that significant attracting forces exist between gas molecules (or atoms, in the case of monoatomic gases), which exhibit a tendency to form a new type of bond. An important landmark in the history of understanding these attracting forces is represented by the liquefaction of helium in experiments by Kamerlingh-Onnes. The very existence of liquid helium provides a most decisive argument about the existence of attractive intermolecular forces acting even between small spherical rare gas atoms such as helium, not bearing any charge or permanent electric multipole moment.

The formation of these special van der Waals bonds, compared to chemical bonds, is not energetically demanding at all; these bonds are, under general laboratory conditions, easily formed and just as easily split. What happens to appear as a weakness represents, surprisingly, a strength of such bonds. In the context of a scenario for the evolution of life on Earth it was necessary to find, besides strong covalent bonding, another type of much weaker bonding allowing easy reversibility of the formation process. The supermolecule formed should allow for repeating opening and closing without changing any important structural feature.

Many years later, in 1930, Fritz London (and soon afterwards Hans Hellmann) made a fundamental step in describing and interpreting these bonds. This was only possible using the recently born quantum mechanics. Though several contributions can be interpreted by classical physics, the most important ones giving rise to the repulsion and attraction between systems (exchange-repulsion and dispersion contributions, see later) that are of quantum origin and could be interpreted only by using the theoretical apparatus of quantum mechanics. Works of these and other pioneers are mentioned or outlined in the classic book on intermolecular interactions by J. O. Hirschfelder, C. F. Curtiss, and R. B. Bird, and a survey of monographs and reviews up to the mid-1980s is presented in a book on intermolecular complexes. Selected summarising works since about 1985 are presented in ref. 7. Specifically to be mentioned are three thematic issues of Chemical Reviews devoted to non-covalent interactions, which appeared in 1988, 1994 and 2000, and one thematic issue of Phys. Chem. Chem. Phys. devoted to the same subject in 2008; all these thematic issues were edited by one of us (PH). Besides these works three books need to be mentioned that supplement the present book. The first one by A. J. Stone focuses on the theory of non- covalent interactions and perturbation calculations of the cluster interaction energy. The second one by A. Karshikoff describes non-covalent interactions in proteins. The third one by I. G. Kaplan deals with the theory and computation of intermolecular interactions. The book presented here is largely based on our theoretical and experimental papers published in the last two decades, which are cited at the end of each chapter. A special place is held by our recent review entitled, "The World of Non-covalent Interactions: 2006" by both present authors and Rudolf Zahradnik.

1.2 A Remark on Nomenclature of Molecular Complexes

Why are molecular complexes, or molecular clusters, as they are most often called, of such interest? The main feature of molecular clusters is that they can be prepared experimentally in supersonic jet expansions and molecular beams as isolated systems exhibiting intermolecular bonds that originate from non-covalent interactions. From the theoretical point of view molecular clusters can also be studied using standard ab initio quantum-chemical methods, treating the cluster as a "supermolecule" composed of several moieties held together by non-covalent bonds.

An issue in the literature that sometimes is unclear relates to the definition of non-covalently bound complexes. A significant feature of such complexes is that the subsystems, of which they are constituted, are not bound by covalent interactions but solely by electric multipole-electric multipole interactions. We consider, however, not only permanent, but also inductive, and time-dependent multipoles. While it is possible to ascribe the stability of a complex to a bond, non-covalent in nature, it is not always easy to localise such a bond in space. When possible it is highly desirable to use another symbol for this bond than that which represents a covalent bond, i.e. a short full line – hence, three dots ... may serve as a representation of a non-covalent bond. The hydrogen molecule and the helium (van der Waals) molecule are adequate representatives: H–H and He ... He, or alternatively H2 and (He)2.

The second type of bond illustrated above still does not have a definite name. No doubt, it is possible to call it a non-covalent bond. Another label, which is sometimes used, is derived from the term weak interactions and therefore the name "weak bonds" is used. This is an unfortunate name, because it is derived from a designation that has been used for a long time in physics in a completely different context. We have favoured for years the designation "van der Waals" (vdW in abbreviated form), e.g., vdW interactions, forces, bonds. It is unfortunately true that this designation has been corrupted – sometimes by poorly defined use – for a component of the empirical force field. In the case of empirical potentials the vdW term means a sum of London dispersion and exchange-repulsion terms. In our previous review we decided to use the term "non-covalent" to classify interactions that are not covalent. We are aware that this definition is again not straightforward and unambiguous since, for example, metallic interactions are also covered but we believe that the term non-covalent properly describes the origin and nature of these interactions. In the very broadest sense non-covalent interactions include electrostatic interactions between permanent multipoles (charge-charge, charge-dipole, charge-multipole, multipole-multi-pole ...), induction and/or polarisation interactions between permanent and induced multipoles, dispersion interactions between instantaneous and induced multipoles and also charge-transfer, ionic and metallic interactions, and interactions leading to formation of H-bonding, halogen bonding and lithium bonding.

1.3 Purpose and Scope: Theory and Experiment

A purpose of this textbook is to illustrate why non-covalent interactions are of fundamental importance for chemistry and why their understanding is a conditio sine qua non for the molecular biodisciplines. Moreover, an attempt will be made to describe correct procedures for treating these interactions theoretically. Those who deal with this subject daily do not need such a recommendation. However, nowadays it is increasingly common for chemists to do the necessary calculations themselves and to develop experience in this field. In contrast to the realm where only chemical (i.e. covalent) bonds play a role, in the area of non-covalently bound complexes it is, in general, not trivial to assert how to proceed and which method and what level of theory guarantees obtaining reliable results. The situation is even more involved because problems of practical value in chemistry and still more in biology are – with respect to computer size – rather extensive. The choice of an appropriate method is especially challenging in these instances. Here, we will describe the main computational procedures to obtain static characteristics of non-covalent species; those characteristics are essential for the understanding of their dynamics.

1.4 Covalent Versus Non-covalent Bonds

The concept of covalent bonding belongs to the most successful concepts in modern science and is, at a certain level, a more or less closed chapter. After about eighty years of intense study, the processes of formation and breaking of covalent bonds are well understood and reliable descriptions of these processes can be performed at various theoretical levels. Calculated characteristic molecular properties agree well with the relevant experiments and there are no fundamental disagreements between state-of-the art theory and experiment. In contrast, the understanding of the nature of non-covalent interactions is far less clear and the respective calculations yield results that are frequently in conflict with experimental data. The basic principles of non-covalent interactions, for instance the hydrogen-bond (H-bond), by Linus Pauling, were formulated in the 1930s. However, despite enormous progress made in theory as well as in experimental techniques in the last decades, we are still far from obtaining unambiguous and quantitatively satisfactory information about non-covalent complexes. Experiments do not yet yield full information on a complex being studied and combining various techniques introduces some ambiguity. Theory, on the other hand, is principally capable of providing full information about a non-covalent complex. For example, we can generate basic information such as structure and stabilisation energy and from the knowledge of the wavefunction we can obtain any other desired property.

A covalent bond is formed when two subsystems with unfilled electronic shells start to overlap. At that point, the electron density between them increases and a bond is created. (More specifically, electron density increases in the bonding region and this increase leads to strengthening of the bond. On the other hand, the increase of electron density in the antibonding region results in weakening of the bond.) The most efficient overlap arises at interatomic distances below 2 Å and at distances larger than about 4 Å overlap is negligible. Non-covalent interactions are, however, known to exist at much greater distances, sometimes at more than 10 Å and in the case of biomacromolecules even at more than 100 Å, which points to the existence of some other source of attraction. The only possibilities are the electric, and to a lesser extent magnetic, properties of the systems. In particular, permanent, induced and instantaneous electric multipoles are sources of attraction (or repulsion) between systems of different types. Electrostatic interaction, using the usual terminology, applies to subsystems with permanent multipole moments each. For two subsystems this interaction energy between charges (= monopoles), dipoles, quadrupoles and higher multipoles is proportional to the product of the multipoles and the first or a higher power of reciprocal distance. In most cases electrostatic interaction dominates compared to other energy terms. Induction (polarisation) interaction represents the interaction between permanent and induced multipoles. For instance, if one subsystem has no permanent multipole moments, i.e. it is neutral and spherically symmetric (i.e. an atom), the electrostatic interaction term is zero and induction is responsible for attraction. If one system is of near central symmetry, such as SF6 (whose lowest permanent moment is the hexadecapole) or CH4 (whose lowest permanent moment is the octopole) electrostatic interaction and induction may be of comparable magnitude. Systems with vanishing permanent moments, for instance, two nonpolar and spherically symmetric systems, exhibit attraction as well. This experimental finding manifested by the liquefaction of noble gases in the early 20th century was at that time very surprising. Only on the basis of quantum mechanics was it possible to theoretically derive the attraction between noble-gas atoms in terms of the London dispersion energy. At that time this explanation represented an important success of the recently born quantum mechanics. London's theory took into account the oscillations of electron clouds and nuclei leading to the generation of instantaneous multipoles, which vanish when integrated over time. For this case one talks about instantaneous multipole-induced multipole moment interactions. The dispersion energy, which is of quantum origin, is proportional to the product of the polarisabilities of the subsystems and a sixth (or higher) power of reciprocal distance. It was believed for a long time that dispersion energy was always smaller than the two previously mentioned energy contributions and that it stabilised mainly noble gas atoms. It has now been shown that dispersion energy between aromatic systems with delocalised electrons is substantial, and stabilisation of these structures is comparable to that of other, for example, hydrogen-bonded structures. This finding has cast new light on the nature of the stabilisation of such biomacromolecules as DNA and proteins. All three energy contributions (electrostatic, induction, dispersion) can be of attractive nature and as such should be balanced by some repulsive force. The so-called exchange-repulsion energy is of quantum-chemical nature too and becomes important when two subsystems overlap. Unlike for the covalent bond, where the electron density between subsystems having unpaired electrons increases in the bonding region, here (subsystems with occupied electron shells) the electron density increases in an antibonding region, which results in mutual repulsion. Most non-covalent complexes exhibit nonzero overlap of their electron clouds. Consequently, some charge transfer between both subsystems exists, giving rise to a small covalent contribution. This is covered by the charge-transfer energy contribution realised not only between the electron donor and the electron acceptor (e.g. benzene ... tetracyanoethylene), but also, more generally, between the proton donor and the proton acceptor (e.g. hydrogen bonding). Only a very small portion of the electron ([approximately equal to] 0.01) is transferred between the subsystems. It must be kept in mind that even chemical processes are connected with only slight changes of the total density and that in the case of non-covalent interactions these changes should be even smaller.

Stabilisation of all the mentioned types of non-covalent complexes is due to favorable energy. This means that the energy of a complex is lower than the sum of the energies of its separated subsystems, which occurs systematically if a complex is formed in vacuo. The situation in an environment and mainly in the water phase is different. This is due to the fact that equilibrium of any non-covalent (as well as covalent) process is not determined by the change of energy (enthalpy) but by the change of the Gibbs energy (ΔG=Δ H – TΔS). The binding can be now realised not only because of favorable energy (enthalpy) but also because of favorable entropy (ΔS is positive, i.e. entropy at the right side of a reaction is higher than that at the left side); in this case the enthalpy (energy) change can be even unfavorable (ΔH>0). This is the so-called hydrophobic interaction, for which the term "hydrophobic bond" is sometimes (albeit incorrectly) used. Evidently, the origin of the stabilisation is completely different. Let us again mention that hydrophobic interactions never occurred for systems in vacuo and is mostly connected with the water environment.

Comparison of theoretical and experimental results is of vital importance for theory as well as for experiments because it allows for the testing of the ability and accuracy of newly developed procedures and techniques. The combination of experiment and theory also gives a deeper insight into the problem studied and so leads to a deeper understanding.

1.5 Experimental Observables

The first question is which properties of non-covalent complexes are observable explicitly? The surprising answer is not so many of them! The structure is not directly observable and can only be determined by measuring the rotational constants thus providing the three principal moments of inertia. Rotational constants, however, do not provide an unambiguous answer concerning structure and geometry (see Section 2.1). A similar situation exists for the determination of stabilisation energies and of the various experimental techniques available only zero-electron kinetic energy (ZEKE) spectroscopy provides directly measurable high-accuracy information on stabilisation energies (see Section 2.3). In addition, directly observable characteristics of a non-covalent complex are vibrational frequencies, not all of which may be seen due to Franck–Condon factors or symmetry selection rules.

(Continues...)

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Edition details

• Publisher: Royal Society of Chemistry
• Publication date: 2009
• Language: English
• ISBN 10: 1847558534
• ISBN 13: 9781847558534
• Binding: Hardcover
• Number of pages: 238