Preface

The chemical bond is the backbone of chemistry. It defines chemistry as the science of understanding and transformation of the physical world in terms of interatomic interactions, which are considered as chemical bonds. Thus, chemistry appears right at the beginning as a fuzzy discipline because the distinction between chemical and nonchemical interatomic interactions is not exactly defined, which creates ongoing controversial debates. But the fuzziness of chemical bonding posed no obstacle for the advancement of chemical research from an esoteric playground to a highly sophisticated academic discipline, which became the basis for industrial growth and economic wealth. In the absence of a physical understanding of chemical bonding, chemists used their imagination and creativity for designing models that proved to be very useful for rationalizing experimental observations.

The development of chemical bonding models, which is an integral part of the progress in experimental chemistry, is a fascinating chapter in the history of mankind. It goes beyond the mere realm of natural science and is part of the evolution of human culture. Chemists analyzed and synthesized in the past centuries a steadily increasing number of new compounds, which required a systematic ordering system to become comprehensible. In order to understand the enormous diversity of molecules and solids, which constitute the chemical universe, chemists developed bonding models that served two purposes. One purpose was to provide an understanding for the observed species, which were classified according to well-defined rules. The second purpose was to establish a guideline for new experiments; a goal that needed a scientific hypothesis in order to distinguish research from random activity.

Originally, chemists suggested bonding models that appeared as simple sketches of connections between atoms that were finally recognized as the elementary building blocks of molecules and solids. The heuristically developed models were continuously refined like a neural network where experimental observations and hypotheses served as guidelines for improving the patterns and archetypes that were used to rationalize new findings. Already in the nineteenth century, chemists realized that atoms might possess different valencies for the formation of chemical bonds without knowing their physical meaning. It is amazing how much progress was made in experimental chemical research without actually knowing much about the constitution of atoms and the nature of the interatomic interactions. The bonding models became the language by which chemists circulate information about molecular structures and reactivities using simple terms that were refined along with the progress in chemical research.

An important step for bridging the gap between chemistry and physics was made in 1916 when G.N. Lewis suggested that the bond line which so far was only used as pictorial representation for a chemical bond without physical meaning, should be identified with a pair of electrons. Lewis knew that his suggestion did not explain the nature of the chemical bond in terms of basic laws of physics, because the classical expression for electrostatic interactions did not agree at all with experimental data. Lewis proposed his model despite the defiance of the physical laws, because of his firm belief that the immense body of chemical facts supported his idea (e.g., the preponderance of molecules with an even number of electrons). Thus, to justify his belief, he postulated that in the atomic world there might be different forces than in the macroscopic world. This intuitively derived model of electron-pair bonding is still the most commonly used archetype for a chemical bond. He could not have foreseen the revolution, which followed from quantum theory that was introduced by Schrödinger and Heisenberg in 1926. One year later, Heitler and London published their landmark paper where they applied quantum theory to describe the interactions between two hydrogen atoms in the bonding and antibonding state of H2. It was the birth of quantum chemistry, which provided the first physically sound description of the covalent chemical bond.

It is remarkable that the work by Heitler and London that outlined for the first time a physically correct description of the chemical bond did not replace the Lewis picture of electron-pair bonding that was based on intuition rather than on elementary physics. One reason is the dramatically different appeal of the two approaches for human imagination of the chemical bond. The Lewis picture is simple to use and it proved as extremely powerful ordering scheme for molecular structures and reactivities. Chemists are generally happy with such models. The quantum theoretical description of interatomic interactions introduced the wave function Ψ as the central term for chemical bonding, which is in contrast an elusive object for human imagination, as evidenced by the intensive discussions about the meaning and the interpretation of Ψ mainly in the physics community.

An important step for building a bridge between the Lewis picture of chemical bonding and quantum theory was made by Pauling in his seminal book The Nature of the Chemical Bond, which was published in 1939. Pauling's work showed that the heuristically derived electron-pair paradigm of Lewis could be dressed by quantum theory keeping the notion of localized bonds in terms of the valence bond (VB) approach. It is thus not surprising that the VB model was well accepted by the chemical community, which quickly adapted the VB notions such as resonance and hybridization for discussing molecular structures and reactivities. The alternative approach of molecular orbital (MO) theory that was developed by Mulliken and Hund was initially met with scepticism by most chemists, because the picture of a localized chemical bond did not seem to be contained in the delocalized MOs. The resistance against MO theory did not change by the fact that phenomena such as the stability of aromatic compounds and spectroscopic data could easily be explained with MOs.

The situation gradually changed during the 1950s till the 1970s, when the advantages of implementing the MO theory became more and more apparent. The development of computer codes by Dewar and Pople, first for semiempirical approaches and then for ab initio methods, paved the way for the acceptance of MO methods. MO theory could much easier become coded into computer programs, which along with the dramatic development of computer hardware produced numerical results with increasing accuracy. Ruedenberg showed that the delocalized MOs could be converted into localized orbitals via unitary transformations, which recovered the Lewis picture even from MO calculations. The breakthrough came with the big success of MO theory of not only explaining but also predicting the reaction course and the stereochemistry of pericyclic reactions. The MO-based frontier orbital theory and Woodward–Hoffmann rules became a standard model for chemical bonding and reactivity, which culminated in awarding the 1981 Nobel Prize to Kenichi Fukui and Roald Hoffmann.

While VB theory nearly disappeared during that time from the horizon of quantum chemistry, it remained remarkably alive in the description of molecular structures and bonding by experimental chemists. In spite of the sweeping success of MO-based quantum chemical methods, which were increasingly used by the general chemical community, the qualitative models that chemists continued to use for sketching molecules and chemical reactions rested often on the picture of localized two-electron bonds. This is because of the unsurpassed simplicity and usefulness of the Lewis bonding model and the associated rules. What can be observed in 2014 is a complementary coexistence of VB and MO models, where the choice of a chemist for answering a question depends on the particular problem and his preference for a specific approach. Also, there has been a remarkable renaissance of VB methods in recent years, which provide an arsenal of bonding models that have been very helpful for explaining molecular structures and reactivities.

The development of quantum chemical methods focussed for a long time on more accurate techniques and efficient algorithms for obtaining numerical results for increasingly larger molecules and for the calculation of reaction pathways. The famous request by Charles Coulson Give us insight, not numbers seemed to have been buried under the quest for more reliable data and little attention was paid to the interpretation of the calculated numbers. The situation has clearly changed in the past two decades. Numerous methods were developed, which aimed at building a bridge between the wealth of numerical data and conceptual models that convert the calculated results into a qualitative understanding of molecular structures and reactivities. Quantum theoretical results can often be presented in terms of figures and schemes rather than merely by presenting tables with numbers, which appeals to the aptitude of most experimental chemists for sensory perception. This does not mean that traditional models such as the Lewis electron pair for chemical bonding have to be abandoned. On the contrary, well-defined partitioning schemes have been introduced, which make it possible to assign calculated numbers that come from accurate quantum chemical calculations to classical concepts replacing handwaving arguments. Although such numbers are not measurable values, they can often be interpreted in terms of physically meaningful expressions, which provide a well-defined ordering scheme for molecular structures and reactivities.

Volume 1 of this book has 11 chapters that present and discuss the...