Showing papers on "Chemical bond published in 2007"

The Quantum theory of atoms in molecules : from solid state to DNA and drug design

Abstract: Foreword. Preface. List of Abbreviations Appearing in this Volume. List of Contributors. 1 An Introduction to the Quantum Theory of Atoms in Molecules (Cherif F. Matta and Russell J. Boyd). 1.1 Introduction. 1.2 The Topology of the Electron Density. 1.3 The Topology of the Electron Density Dictates the Form of Atoms in Molecules. 1.4 The Bond and Virial Paths, and the Molecular and Virial Graphs. 1.5 The Atomic Partitioning of Molecular Properties. 1.6 The Nodal Surface in the Laplacian as the Reactive Surface of a Molecule. 1.7 Bond Properties. 1.8 Atomic Properties. 1.9 "Practical" Uses and Utility of QTAIM Bond and Atomic Properties. 1.10 Steps of a Typical QTAIM Calculation. References. Part I Advances in Theory. 2 The Lagrangian Approach to Chemistry (Richard F. W. Bader). 2.1 Introduction. 2.2 The Lagrangian Approach. 2.3 The Action Principle in Quantum Mechanics. 2.4 From Schrodinger to Schwinger. 2.5 Molecular Structure and Structural Stability. 2.6 Re.ections and the Future. References. 3 Atomic Response Properties (Todd A. Keith). 3.1 Introduction. 3.2 Apparent Origin-dependence of Some Atomic Response Properties. 3.3 Bond Contributions to "Null" Molecular Properties. 3.4 Bond Contributions to Atomic Charges in Neutral Molecules. 3.5 Atomic Contributions to Electric Dipole Moments of Neutral Molecules. 3.6 Atomic Contributions to Electric Polarizabilities. 3.7 Atomic Contributions to Vibrational Infrared Absorption Intensities. 3.8 Atomic Nuclear Virial Energies. 3.9 Atomic Contributions to Induced Electronic Magnetic Dipole Moments. 3.10 Atomic Contributions to Magnetizabilities of Closed-Shell Molecules. References. 4 QTAIM Analysis of Raman Scattering Intensities: Insights into the Relationship Between Molecular Structure and Electronic Charge Flow (Kathleen M. Gough, Richard Dawes, Jason R. Dwyer, and Tammy L. Welshman). 4.1 Introduction. 4.2 Background to the Problem. 4.3 Methodology. 4.4 Speci.c Examples of the Use of AIM2000 Software to Analyze Raman Intensities. 4.5 Patterns in alpha That Are Discovered Through QTAIM. 4.6 Patterns in qa/qrCH That Apply Across Di.erent Structures, Conformations, Molecular Types: What is Transferable? 4.7 What Can We Deduce From Simple Inspection of &delta &alpha /&delta rCH and &delta &alpha /&delta rCC From Gaussian? 4.8 Conclusion. References. 5 Topological Atom-Atom Partitioning of Molecular Exchange Energy and its Multipolar Convergence (Michel Rafat and Paul L. A. Popelier). 5.1 Introduction. 5.2 Theoretical Background. 5.3 Details of Calculations. 5.4 Results and Discussion. 5.5 Conclusion. References. 6 The ELF Topological Analysis Contribution to Conceptual Chemistry and Phenomenological Models (Bernard Silvi and Ronald J. Gillespie). 6.1 Introduction. 6.2 Why ELF and What is ELF? 6.3 Concepts from the ELF Topology. 6.4 VSEPR Electron Domains and the Volume of ELF Basins. 6.5 Examples of the Correspondence Between ELF Basins and the Domains of the VSEPR Model. 6.6 Conclusions. References. Part II Solid State and Surfaces. 7 Solid State Applications of QTAIM and the Source Function - Molecular Crystals, Surfaces, Host-Guest Systems and Molecular Complexes (Carlo Gatti). 7.1 Introduction. 7.2 QTAIM Applied to Solids - the TOPOND Package. 7.3 QTAIM Applied to Molecular Crystals. 7.4 QTAIM Applied to Surfaces. 7.5 QTAIM Applied to Host-Guest Systems. 7.6 The Source Function: Theory. References. 8 Topology and Properties of the Electron Density in Solids (Victor Luana, Miguel A. Blanco, Aurora Costales, Paula Mori-Sanchez, and Angel Martin Pendas). 8.1 Introduction. 8.2 The Electron Density Topology and the Atomic Basin Shape. 8.3 Crystalline Isostructural Families and Topological Polymorphism. 8.4 Topological Classi.cation of Crystals. 8.5 Bond Properties - Continuity from the Molecular to the Crystalline Regime. 8.6 Basin Partition of the Thermodynamic Properties. 8.7 Obtaining the Electron Density of Crystals. References. 9 Atoms in Molecules Theory for Exploring the Nature of the Active Sites on Surfaces (Yosslen Aray, Jesus Rodriguez, and David Vega). 9.1 Introduction. 9.2 Implementing the Determination of the Topological Properties of p(r) from a Three-dimensional Grid. 9.3 An Application to Nanocatalyts - Exploring the Structure of the Hydrodesulfurization MoS2 Catalysts. References. Part III Experimental Electron Densities and Biological Molecules. 10 Interpretation of Experimental Electron Densities by Combination of the QTAMC and DFT (Vladimir G. Tsirelson). 10.1 Introduction. 10.2 Specificity of the Experimental Electron Density. 10.3 Approximate Electronic Energy Densities. 10.4 The Integrated Energy Quantities. 10.5 Concluding Remarks. References. 11 Topological Analysis of Proteins as Derived from Medium and Highresolution Electron Density: Applications to Electrostatic Properties (Laurence Leherte, Benoit Guillot, Daniel P. Vercauteren, Virginie Pichon-Pesme, Christian Jelsch, Ange'lique Lagoutte, and Claude Lecomte). 11.1 Introduction. 11.2 Methodology and Technical Details. 11.3 Topological Properties of Multipolar Electron Density Database. 11.4 Analysis of Local Maxima in Experimental and Promolecular Mediumresolution Electron Density Distributions. 11.5 Calculation of Electrostatic Properties from Atomic and Fragment Representations of Human Aldose Reductase. 11.6 Conclusions and Perspectives. References. 12 Fragment Transferability Studied Theoretically and Experimentally with QTAIM - Implications for Electron Density and Invariom Modeling (Peter Luger and Birger Dittrich). 12.1 Introduction. 12.2 Experimental Electron-density Studies. 12.3 Studying Transferability with QTAIM - Atomic and Bond Topological Properties of Amino Acids and Oligopeptides. 12.4 Invariom Modeling. 12.5 Applications of Aspherical Invariom Scattering Factors. 12.6 Conclusion. References. Part IV Chemical Bonding and Reactivity. 13 Interactions Involving Metals - From "Chemical Categories" to QTAIM, and Backwards (Piero Macchi and Angelo Sironi). 13.1 Introduction. 13.2 The Electron Density in Isolated Metal Atoms - Hints of Anomalies. 13.3 Two-center Bonding. 13.4 Three-center Bonding. 13.5 Concluding Remarks. References. 14 Applications of the Quantum Theory of Atoms in Molecules in Organic Chemistry - Charge Distribution, Conformational Analysis and Molecular Interactions (Jesus Hernandez-Trujillo, Fernando Cortes-Guzman, and Gabriel Cuevas). 14.1 Introduction. 14.2 Electron Delocalization. 14.3 Conformational Equilibria. 14.4 Aromatic Molecules. References. 15 Aromaticity Analysis by Means of the Quantum Theory of Atoms in Molecules (Eduard Matito, Jordi Poater, and Miquel Sola ). 15.1 Introduction. 15.2 The Fermi Hole and the Delocalization Index. 15.3 Electron Delocalization in Aromatic Systems. 15.4 Aromaticity Electronic Criteria Based on QTAIM. 15.5 Applications of QTAIM to Aromaticity Analysis. 15.6 Conclusions. References. 16 Topological Properties of the Electron Distribution in Hydrogen-bonded Systems (Ignasi Mata, Ibon Alkorta, Enrique Espinosa, Elies Molins, and Jose Elguero). 16.1 Introduction. 16.2 Topological Properties of the Hydrogen Bond. 16.3 Energy Properties at the Bond Critical Point (BCP). 16.4 Topological Properties and Interaction Energy. 16.5 Electron Localization Function, n(r). 16.6 Complete Interaction Range. 16.7 Concluding Remarks. References. 17 Relationships between QTAIM and the Decomposition of the Interaction Energy - Comparison of Different Kinds of Hydrogen Bond (Stawomir J. Grabowski). 17.1 Introduction. 17.2 Diversity of Hydrogen-bonding Interactions. 17.3 The Decomposition of the Interaction Energy. 17.4 Relationships between the Topological and Energy Properties of Hydrogen Bonds. 17.5 Various Other Interactions Related to Hydrogen Bonds. 17.6 Summary. References. Part V Application to Biological Sciences and Drug Design. 18 QTAIM in Drug Discovery and Protein Modeling (Nagamani Sukumar and Curt M. Breneman). 18.1 QSAR and Drug Discovery. 18.2 Electron Density as the Basic Variable. 18.3 Atom Typing Scheme and Generation of the Transferable Atom Equivalent (TAE) Library. 18.4 TAE Reconstruction and Descriptor Generation. 18.5 QTAIM-based Descriptors. 18.6 Sample Applications. 18.7 Conclusions. References. 19 Fleshing-out Pharmacophores with Volume Rendering of the Laplacian of the Charge Density and Hyperwall Visualization Technology (Preston J. MacDougall and Christopher E. Henze). 19.1 Introduction. 19.2 Computational and Visualization Methods. 19.3 Subatomic Pharmacophore Insights. 19.4 Conclusion. References. Index.

Red-, blue-, or no-shift in hydrogen bonds: a unified explanation.

Abstract: We provide a simple explanation for X-H bond contraction and the associated blue shift and decrease of intensity in IR spectrum of the so-called improper hydrogen bonds This explanation organizes hydrogen bonds (HBs) with a seemingly random relationship between the X-H bond length (and IR frequency and its intensity) to its interaction energy The factors which affect the X-H bond in all X-H [midline ellipsis] Y HBs can be divided into two parts: (a) The electron affinity of X causes a net gain of electron density at the X-H bond region in the presence of Y and encourages an X-H bond contraction (b) The well understood attractive interaction between the positive H and electron rich Y forces an X-H bond elongation For electron rich, highly polar X-H bonds (proper HB donors) the latter almost always dominates and results in X-H bond elongation, whereas for less polar, electron poor X-H bonds (pro-improper HB donors) the effect of the former is noticeable if Y is not a very strong HB acceptor Although both the above factors increase with increasing HB acceptor ability of Y, the shortening effect dominates over a range of Ys for suitable pro-improper X-Hs resulting in a surprising trend of decreasing X-H bond length with increasing HB acceptor ability The observed frequency and intensity variations follow naturally The possibility of HBs which do not show any IR frequency change in the X-H stretching mode also directly follows from this explanation

Hydrogen multicentre bonds

Abstract: The concept of a chemical bond stands out as a major development in the process of understanding how atoms are held together in molecules and solids. Lewis’ classical picture of chemical bonds as shared-electron pairs1 evolved to the quantum-mechanical valence-bond and molecular-orbital theories2,3, and the classification of molecules and solids in terms of their bonding type: covalent, ionic, van der Waals and metallic. Along with the more complex hydrogen bonds4 and three-centre bonds5,6, they form a paradigm within which the structure of almost all molecules and solids can be understood. Here, we present evidence for hydrogen multicentre bonds—a generalization of three-centre bonds—in which a hydrogen atom equally bonds to four or more other atoms. When substituting for oxygen in metal oxides, hydrogen bonds equally to all the surrounding metal atoms, becoming fourfold coordinated in ZnO, and sixfold coordinated in MgO. These multicentre bonds are remarkably strong despite their large hydrogen–metal distances. The calculated local vibration mode frequency in MgO agrees with infrared spectroscopy measurements7. Multicoordinated hydrogen also explains the dependence of electrical conductivity on oxygen partial pressure, resolving a long-standing controversy on the role of point defects in unintentional n-type conductivity of ZnO (refs 8–10).

Structure of phase III of solid hydrogen

Abstract: Hydrogen, being the first element in the periodic table, has the simplest electronic structure of any atom, and the hydrogen molecule contains the simplest covalent chemical bond. Nevertheless, the phase diagram of hydrogen is poorly understood. Determining the stable structures of solid hydrogen is a tremendous experimental challenge1,2,3, because hydrogen atoms scatter X-rays only weakly, leading to low-resolution diffraction patterns. Theoretical studies encounter major difficulties owing to the small energy differences between structures and the importance of the zero-point motion of the protons. We have systematically investigated the zero-temperature phase diagram of solid hydrogen using first-principles density functional theory (DFT) electronic-structure methods4, including the proton zero-point motion at the harmonic level. Our study leads to a radical revision of the DFT phase diagram of hydrogen up to nearly 400 GPa. That the most stable phases remain insulating to very high pressures eliminates a major discrepancy between theory5 and experiment6. One of our new phases is calculated to be stable over a wide range of pressures, and its vibrational properties agree with the available experimental data for phase III.

Orbital Reconstruction and Covalent Bonding at an Oxide Interface

Abstract: Orbital reconstructions and covalent bonding must be considered as important factors in the rational design of oxide heterostructures with engineered physical properties. We have investigated the interface between high-temperature superconducting (Y,Ca)Ba(2)Cu3O7 and metallic La(0.67)Ca(0.33)MnO3 by resonant x-ray spectroscopy. A charge of about -0.2 electron is transferred from Mn to Cu ions across the interface and induces a major reconstruction of the orbital occupation and orbital symmetry in the interfacial CuO2 layers. In particular, the Cu d(3z(2)-r(2)) orbital, which is fully occupied and electronically inactive in the bulk, is partially occupied at the interface. Supported by exact-diagonalization calculations, these data indicate the formation of a strong chemical bond between Cu and Mn atoms across the interface. Orbital reconstructions and associated covalent bonding are thus important factors in determining the physical properties of oxide heterostructures.

Hydrogen bond dynamics in aqueous NaBr solutions

Abstract: hydrogen bond dynamics in aqueous solutions (20) and other hydrogen-bonded liquids (21). Water molecules around monatomic ions provide the simplest system for the investigation of the dynamics of water in the presence of charges. Therefore, examining the dynamics of water in simple ionic solutions, such as sodium bromide (NaBr), can provide valuable insights that will increase understanding of chemical processes and biological systems that involve charged species in aqueous solutions. In pure water, water molecules are hydrogen-bonded to neighboring water molecules in a more or less tetrahedral geometry making an extended hydrogen bond network. Hydrogen bonds are continually breaking and reforming, and hydrogen bond lengths (strengths) are continually changing (9). The structure of water fluctuates on femtosecond to picosecond time scales (6, 9). The slowest component of the fluctuations is associated with the global structural rearrangement of the hydrogen bond network (7). As salt is dissolved in water, water molecules form hydration shells around ions, and consequently the local hydrogen bond network is perturbed. In this paper, a very detailed study of the hydrogen bond dynamics of water in highly concentrated NaBr solutions is presented. Four experimental observables, which are based on measurements on the OD stretch mode of HOD in H2O, were used. Ultrafast 2D IR vibrational echo experiments were used to measure spectral diffusion, which is directly related to the time evolution of the hydrogen bond structure. Polarization-selective IR pump–probe experiments were used to measure orientational relaxation dynamics of water, which is affected by the interactions of water molecules with their surroundings. The IR pump–

Bond paths as privileged exchange channels.

Abstract: Evidence that the bond paths of the quantum theory of atoms-in-molecules (QTAIM) signal preferred quantum-mechanical exchange channels is presented. We show how bond paths between an atom A and the atoms B in its environment appear to be determined by competition among the A-B exchange-correlation energies that always contribute to stabilize the A-B interactions. These pairwise additive stabilizations depend neither on the attractive or repulsive nature of the classical electrostatic interaction between the atoms' charge densities, nor on the change in the self energies of the atoms involved. These other terms may well cause an overall molecular-energy increase in spite of a possibly large A-B exchange-correlation stabilization. After our proposal, bond paths, both at and out of equilibrium geometries, are endowed with a specific energetic meaning that should contribute to reconcile the orthodox QTAIM interpretation with other widely accepted views, and to settle recent controversies questioning the meaning of hydrogen-hydrogen bonding and the nature of the so-called "steric interactions", the role of bond paths in endohedral complexes, and the generality of the results provided by the QTAIM. Implications for the nature of more general closed-shell interactions are also briefly discussed.

Reaching the maximum multiplicity of the covalent chemical bond.

Abstract: The bond order and in particular the possibility of multiple bonding between atoms in a molecule have been highlighted in two recent articles. 2] Theoretical and experimental work have challenged old chemical paradigms concerning the possible multiplicity that can be achieved in a chemical bond. On the other hand, the concept of a multiple bond is not clearly defined and there is a need for a more quantitative measure. In this contribution we attempt to introduce such a measure and apply it to a number of multiply bonded systems. As a result of the analysis, we show that the highest multiplicity that can be achieved in a bond between two equal atoms is six. The multiplicity of a chemical bond is determined by the number of electron pairs that occupy the region between the two bonded atoms in bonding molecular orbitals. The hydrogen molecule has, for example, a single bond with two electrons in one orbital formed from the 1s orbitals on each atom. The nitrogen molecule, N2, has a triple bond; the three unpaired 2p electrons on each atom combine to form this very strong bond. Before 1964, the triple bond was assumed to be the highest multiplicity that a chemical bond can have. We show here, through a systematic study of the covalent chemical bond covering the entire periodic system, that the maximum bond multiplicity is six. The maximum value is reached by the tungsten diatom, W2. No other pair of atoms in the periodic system (atomic numbers smaller than about 100) reaches a higher bond order. A single covalent chemical bond between two atoms is, in simple molecular orbital (MO) theory, described by a bonding orbital occupied by two electrons. This is, however, an oversimplified picture of bonding that only works for strong bonds and near the equilibrium geometry. None would say that there is a chemical bond between two hydrogen atoms that are at a distance of 100 . from each other. However, this is the picture that emerges from the simple theory. A more accurate description uses two orbitals to describe the bond: a bonding orbital and the corresponding antibonding orbital. Both orbitals are occupied in the true wave function of the molecule. Let us assume that the occupation of the bonding orbital is hb= 2 x. The occupation of the antibonding orbital will then be close to ha= x, such that the sum is two. When the molecule is close to equilibrium, x will be small for a normal chemical bond, but when the molecule dissociates, x will increase to become one. In this case, there is no chemical bond and the wave function describes a system of two radicals with one electron on each of them.We can use this property of the molecular orbitals to define an effective bond order (EBO) for a single bond as (hb ha)/2, which is then close to one for a normal chemical bond but goes to zero when the bond weakens. For multiply bonded molecules we add up the contributions from each bond to obtain the total EBO. The EBO is non-integer and in naming the multiplicity of a bond one may then use the lowest integer value larger than the EBO. Why is this interesting? If one assumes that the bond formed between the two fragments is weak for some reason, for example, as a result of steric hindrance, then the value of x may be quite different from zero. When the value is 0.5, for instance, the bond is only halfway formed and the effective bond order is 1 x= 0.5. In multiply bonded systems, the different orbitals forming the bonds may have different overlaps and x may vary considerably from bond to bond. This measure of the bondmultiplicity is based on very well defined and stable quantities: the occupation numbers of the natural orbitals (NOs). It can only be used together with wave functions that give realistic values for these quantities. These are by necessity multiconfigurational wave functions. The concept becomes meaningless together with Hartree–Fock or DFTwave functions. It is important to emphasize that the NO occupation numbers are stable quantities that do not vary much when a wave function is improved, once a wave function has been defined that includes the most important NOs. The dependence on the AO basis set is also small, which makes the NOs and their occupation numbers very useful as measures of the bonding in a molecule. As many quantities used to describe what the electrons do in a molecule, the bond order is not a measurable quantity, nor is it directly related to such quantities. Different definitions are therefore possible, from a simple count of electrons to more sophisticated measures based on different partitionings of the density matrix. Such measures are, however, often very methodand basis-set-dependent, which is not the case for the definition applied here. Before 1964, it was assumed that the highest bond order that could exist between two atoms was three. That [*] Prof. B. O. Roos Department of Theoretical Chemistry Chemical Center P.O.B. 124, 22100 Lund (Sweden) Fax: (+46)46-22-4543 E-mail: bjorn.roos@teokem.lu.se

Charge Decomposition Analysis of the Electron Localizability Indicator: A Bridge between the Orbital and Direct Space Representation of the Chemical Bond

Abstract: The novel functional electron localizability indicator is a useful tool for investigating chemical bonding in molecules and solids. In contrast to the traditional electron localization function (ELF), the electron localizability indicator is shown to be exactly decomposable into partial orbital contributions even though it displays at the single-determinantal level of theory the same topology as the ELF. This approach is generally valid for molecules and crystals at either the single-determinantal or the explicitly correlated level of theory. The advantages of the new approach are illustrated for the argon atom, homonuclear dimers N 2 and F 2 , unsaturated hydrocarbons C 2 H 4 and C 6 H 6 , and the transition-metal-containing molecules Sc 2 2+ and TiF 4 .

A Chemist's Guide to Valence Bond Theory

Abstract: Preface. Chapter 1. A Brief Story of Valence Bond Theory, Its Rivalry With Molecular Orbital Theory, Its Demise, And Resurgence. 1.1. Roots of VB Theory. 1.2. Origins of MO Theory and the Roots of VB-MO Rivalry. 1.3. One Theory is Up the Other is Down. 1.4. Mythical Failures of VB Theory: More Ground is Gained by MO Theory. 1.5. Are the Failures of VB Theory Real? 1.6. VB is a Legitimate Theory Alongside Molecular Orbital Theory. 1.7. Modern VB Theory: Valence Bond Theory is Coming of Age. Chapter 2. A Brief Tour Through Some Valence Bond Outputs and Terminology. 2.1. Valence Bond Output for the H 2 Molecule. 2.2. Valence Bond Mixing Diagrams. 2.3. Valence Bond Output for the HF Molecule. Chapter 3. Basic Valence Bond Theory. 3.1. Writing and Representing Valence Bond Wave. 3.2 Overlaps between Determinants. 3.3 Valence Bond Formalism Using the Exact Hamiltonian. 3.4 Valence Bond Formalism using an Effective Hamiltonian. 3.5 Some Simple Formulas for Elementary Interactions. 3.6 Structural Coefficients and Weights of Valence Bond Wave Function. 3.7 Bridges Between Molecular Orbital and Valence Bond Theories. Chapter 4. Mapping Molecular Orbitals-Configuration Interaction to Valence Bond Wave Functions. 4.1. Generating a Set of Valence Bond structures. 4.2. Mapping a Molecular Orbital-Configuration Interaction. 4.3. Using Half-Determinants to Calculate Overlaps between Valence Bond Structures. 5. Are the "Failures" of Valence Bond Theory Real? 5.1. Introduction. 5.2. The Triplet Ground State of Dioxygen. 5.3. Aromaticity-Antiaromaticity in Ionic Rings CnHn+/- 5.4. Aromaticity/Antiaromaticity in Neutral Rings. 5.5. The Valence Ionization Spectrum of CH4 5.6. The Valence Ionization Spectrum of H2O and the "Rabbit-Ear" Lone Pairs. 5.7. A Summary. 6. Valence Bond Diagrams for Chemical Reactivity. 6.1. Introduction. 6.2. Two Archetypal Valence Bond Diagrams. 6.3. The Valence Bond State Correlation Diagram Model and Its General Outlook on Reactivity. 6.4. Construction of Valance Bond State Correlation Diagram Model and Its General Outlook on Reactivity. 6.4. Construction of Valence Bond State Correlation Diagrams for Elementary Processes. 6.5. Barrier Expressions Based on the Valence Bond State Correlation Diagram Model. 6.6. Making Qualitative Reactivity Predictions with the Valence Bond State Correlation Diagram. 6.7. Valence Bond Configuration Mixing Diagrams: General Features. 6.8. Valence Bond Configuration Mixing Diagram with Ionic Intermediate Curves. 6.9. Valence Bond Configuration Mixing Diagram with Intermediates Nascent from "Foreign States". 6.10. Valence Bond State Correlation Diagram: A General Model for Electronic Delocalization in Clusters. 6.11. Valence Bond State Correlation Diagram: Application to Photochemical Reactivity. 6.12. A Summary. 7. Using Valence Bond Theory to Compute and Conceptualize Excited States. 7.1. Excited States of a Single Bond. 7.2. Excited States of Molecules with Conjugated Bonds. 7.3. A Summary. 8. Spin Hamiltonian Valence bond Theory and its Applications to Organic Radicals, Diradicals, and Polyradicals. 8.1. A Topological Semiempirical Hamiltonian. 8.2. Applications. 8.3. A Summary. 9. Currently Available AB Initio Valence Bond Computational Methods and their Principles. 9.1. Introduction. 9.2. Valence Bond Methods Based on Semilocalized Orbitals. 9.3. Valence Bond Methods Based on Localized Orbitals. 9.4. Methods for Getting Valence Bond Quantities. 9.5. A Valence Bond Methods with Polarizable Continuum Model. 10. Do Your Own Valence Bond Calculations-A Practical Guide. 10.1. Introduction. 10.2. Wave Functions and Energies for the Ground State of F2. 10.3. Valence Bond Calculations of Diabatic States and Resonance Energies. 10.4. Comments on Calculations of VBSCDs and VBCMDs. Epilogue. Glossary. Index.

Comprehensive analysis of chemical bonding in boron clusters.

Abstract: We present a comprehensive analysis of chemical bonding in pure boron clusters. It is now established in joint experimental and theoretical studies that pure boron clusters are planar or quasi-planar at least up to twenty atoms. Their planarity or quasi-planarity was usually discussed in terms of pi-delocalization or pi-aromaticity. In the current article, we demonstrated that one cannot ignore sigma-electrons and that the presence of two-center two-electron (2c--2e) peripheral B--B bonds together with the globally delocalized sigma-electrons must be taken into consideration when the shape of pure boron cluster is discussed. The global aromaticity (or global antiaromaticity) can be assigned on the basis of the 4n+2 (or 4n) electron counting rule for either pi- or sigma-electrons in the planar structures. We showed that pure boron clusters could have double (sigma- and pi-) aromaticity (B3-, B4, B5+, B6(2+), B7+, B7-, B8, B(8)2-, B9-, B10, B11+, B12, and B13+), double (sigma- and pi-) antiaromaticity (B6(2-), B15), or conflicting aromaticity (B5-,sigma-antiaromatic and pi-aromatic and B14, sigma-aromatic and pi-antiaromatic). Appropriate geometric fit is also an essential factor, which determines the shape of the most stable structures. In all the boron clusters considered here, the peripheral atoms form planar cycles. Peripheral 2c--2e B--B bonds are built up from s to p hybrid atomic orbitals and this enforces the planarity of the cycle. If the given number of central atoms (1, 2, 3, or 4) can perfectly fit the central cavity then the overall structure is planar. Otherwise, central atoms come out of the plane of the cycle and the overall structure is quasi-planar.

Description of the Ground‐State Covalencies of the Bis(dithiolato) Transition‐Metal Complexes from X‐ray Absorption Spectroscopy and Time‐Dependent Density‐Functional Calculations

Abstract: The electronic structures of [M(L(Bu))(2)](-) (L(Bu)=3,5-di-tert-butyl-1,2-benzenedithiol; M=Ni, Pd, Pt, Cu, Co, Au) complexes and their electrochemically generated oxidized and reduced forms have been investigated by using sulfur K-edge as well as metal K- and L-edge X-ray absorption spectroscopy. The electronic structure content of the sulfur K-edge spectra was determined through detailed comparison of experimental and theoretically calculated spectra. The calculations were based on a new simplified scheme based on quasi-relativistic time-dependent density functional theory (TD-DFT) and proved to be successful in the interpretation of the experimental data. It is shown that dithiolene ligands act as noninnocent ligands that are readily oxidized to the dithiosemiquinonate(-) forms. The extent of electron transfer strongly depends on the effective nuclear charge of the central metal, which in turn is influenced by its formal oxidation state, its position in the periodic table, and scalar relativistic effects for the heavier metals. Thus, the complexes [M(L(Bu))(2)](-) (M=Ni, Pd, Pt) and [Au(L(Bu))(2)] are best described as delocalized class III mixed-valence ligand radicals bound to low-spin d(8) central metal ions while [M(L(Bu))(2)](-) (M=Cu, Au) and [M(L(Bu))(2)](2-) (M=Ni, Pd, Pt) contain completely reduced dithiolato(2-) ligands. The case of [Co(L(Bu))(2)](-) remains ambiguous. On the methodological side, the calculation led to the new result that the transition dipole moment integral is noticeably different for S(1s)-->valence-pi versus S(1s)-->valence-sigma transitions, which is explained on the basis of the differences in radial distortion that accompany chemical bond formation. This is of importance in determining experimental covalencies for complexes with highly covalent metal-sulfur bonds from ligand K-edge absorption spectroscopy.

Unexpected stability of Al4H6: a borane analog?

Abstract: Whereas boron has many hydrides, aluminum has been thought to exhibit relatively few. A combined anion photoelectron and density functional theory computational study of the Al4H –6 anion and its corresponding neutral, Al4H6, showed that Al4H6 can be understood in terms of the Wade-Mingos rules for electron counting, suggesting that it may be a borane analog. The data support an Al4H6 structure with a distorted tetrahedral aluminum atom framework, four terminal Al-H bonds, and two sets of counter-positioned Al-H-Al bridging bonds. The large gap between the highest occupied and the lowest unoccupied molecular orbitals found for Al4H6, together with its exceptionally high heat of combustion, further suggests that Al4H6 may be an important energetic material if it can be prepared in bulk.

Delta aromaticity in [Ta3O3]-.

Abstract: We report experimental and theoretical evidence of -aromaticity, which is discovered in the Ta3O3 – cluster via a combined photoelectron spectroscopy and ab initio study. Well-resolved low-lying electronic transitions are observed in the photoelectron spectra of Ta3O3 – and are compared with ab initio calculations, which show that the Ta3O3 – cluster possesses a planar D3h triangular structure. Chemical bonding analyses reveal that among the five valence molecular orbitals responsible for the multi-center metal-metal bonding there is a completely bonding delta and orbital from the 5d atomic orbitals of Ta. The totally delocalized multi-center bond renders -aromaticity for Ta3O3 – and represents a new mode of chemical bonding. Ta3O3 – is the first -aromatic molecule confirmed experimentally and theoretically, suggesting that -aromaticity may exist in many multi-nuclear, low oxidation state transition-metal compounds.

δ Aromaticity in [Ta3O3]−

Abstract: The concept of aromaticity was introduced into organic chemistry to describe delocalized p bonding in planar, cyclic, and conjugate molecules possessing (4n+2) p electrons. In recent years, this concept has been advanced into main-group molecules including organometallic compounds with cyclic cores of metal atoms and, in particular, all-metal clusters. It has been shown that main-group clusters may exhibit multiple aromaticity (s and p), multiple antiaromaticity (s and p), and conflicting aromaticity (s aromaticity and p antiaromaticity or s antiaromaticity and p aromaticity). Here, we report experimental and theoretical evidence of d aromaticity, which is only possible in transition-metal systems. It is discovered in the [Ta3O3] cluster through combined photoelectron spectroscopy and ab initio studies. Well-resolved low-lying electronic transitions are observed in the photoelectron spectra of [Ta3O3] and are compared with ab initio calculations, which show that the [Ta3O3] cluster has a planar D3h triangular structure. Chemical-bonding analyses reveal that among the five valence molecular orbitals involved in the multicenter metal–metal bonding, there is a completely bonding d and p orbital formed from the 5d atomic orbitals of Ta. The totally delocalized multicenter d bond renders d aromaticity for [Ta3O3] and represents a new mode of chemical bonding. [Ta3O3] is the first d-aromatic molecule confirmed experimentally and theoretically, which suggests that d aromaticity may exist in many multinuclear, lowoxidation-state transition-metal compounds. In 1964, Cotton and co-workers published a milestone work on K2[Re2Cl8]·2H2O, [7] in which they showed the presence of a new type of chemical bond—a d bond between the two Re atoms. Since then, a branch of inorganic chemistry has been developed that involves multiple metal–metal bonding with bond orders higher than three, the maximum allowed for main-group systems. Power and co-workers recently reported the synthesis of a Cr2 compound with a quintuple bond (spd) between the two Cr atoms. This work, along with recent quantum chemical studies of multiple bonds in U2 and [Re2Cl8] 2 , has generated renewed interest in multiple metal–metal bonding. The presence of d bonds between two transition-metal atoms suggests that multicenter transition-metal species with a completely delocalized cyclic d bond may exist, thus raising the possibility of d aromaticity analogous to p or s aromaticity in main-group systems. We have been interested in understanding the electronic structure and chemical bonding of early transition-metal oxide clusters as a function of size and composition, and in using them as potential molecular models for oxide catalysts. During our investigation of tantalum oxide clusters, we found the presence of d aromaticity in the [Ta3O3] cluster, in which each Ta atom is in a low oxidation state of Ta and still possesses three electrons for Ta–Ta bonding. The experiment was conducted by using a magneticbottle-type photoelectron spectroscopy apparatus equipped with a laser vaporization cluster source. [TamOn] clusters with various compositions were produced by laser vaporization of a pure tantalum disk target in the presence of a helium carrier gas seeded with O2, and were size-separated by time-of-flight mass spectrometry. The [Ta3O3] species was mass-selected and decelerated before photodetachment by a pulsed laser beam. Photoelectron spectra were obtained at two relatively high photon energies, 193 nm (6.424 eV) and 157 nm (7.866 eV), to guarantee access to all valence electronic transitions (Figure 1). Three well-resolved bands (X, A, and B) were observed at the lower-binding-energy side. The X band is much more intense and shows a discernible splitting at 193 nm (Figure 1a). Surprisingly, no well-defined electronic transitions were observed beyond 3.7 eV, where continuous signals were present, probably as a result of multielectron transitions. The vertical detachment energies (VDEs) of the observed transitions at the low-bindingenergy side are given in Table 1, where they are compared with theoretical calculations by two different methods. [*] Dr. H. J. Zhai, Prof. Dr. L. S. Wang Department of Physics Washington State University 2710 University Drive, Richland, WA 99354 (USA) and Chemical & Materials Sciences Division Pacific Northwest National Laboratory MS K8–88, P.O. Box 999, Richland, WA 99352 (USA) Fax: (+1)509-376-6066 E-mail: ls.wang@pnl.gov

Structure and bonding of the water-hydroxyl mixed phase on Pt(111)

Abstract: We combine low-energy electron diffraction (LEED), X-ray photoelectron spectroscopy (XPS), X-ray absorption spectroscopy (XAS), and Auger electron spectroscopy (AES) with density functional theory (DFT) to reveal the structure and bonding of water-hydroxyl mixed layers adsorbed on Pt(111). We find that the stable water-hydroxyl adlayer forms a mixed phase of nearly coplanar hexamer structures resulting in (root 3 x root 3)R30 degrees and (3 x 3) unit cells, respectively. In the asymmetric (3 x 3) structure the lateral O-O distances alternate between long and short bond lengths similar to the chemical bonding network for OH- ions in solution. The chemical driving force behind this similarity is discussed in a molecular orbital picture.

Electronic structure of CO--an exercise in modern chemical bonding theory.

Abstract: This paper discusses recent progress that has been made in the understanding of the electronic structure and bonding situation of carbon monoxide which was analyzed using modern quantum chemical methods. The new results are compared with standard models of chemical bonding. The electronic charge distribution and the dipole moment, the nature of the HOMO and the bond dissociation energy are discussed in detail. © 2006 Wiley Periodicals, Inc. J Comput Chem, 2006

Electronic structure of copper phthalocyanine: An experimental and theoretical study of occupied and unoccupied levels

Abstract: An experimental and theoretical study of the electronic structure of copper phthalocyanine (CuPc) molecule is presented. We performed x-ray photoemission spectroscopy (XPS) and photoabsorption [x-ray absorption near-edge structure (XANES)] gas phase experiments and we compared the results with self-consistent field, density functional theory (DFT), and static-exchange theoretical calculations. In addition, ultraviolet photoelectron spectra (UPS) allowed disentangling several outer molecular orbitals. A detailed study of the two highest occupied orbitals (having a(1u) and b(1g) symmetries) is presented: the high energy resolution available for UPS measurements allowed resolving an extra feature assigned to vibrational stretching in the pyrrole rings. This observation, together with the computed DFT electron density distributions of the outer valence orbitals, suggests that the a(1u) orbital (the highest occupied molecular orbital) is mainly localized on the carbon atoms of pyrrole rings and it is doubly occupied, while the b(1g) orbital, singly occupied, is mainly localized on the Cu atom. Ab initio calculations of XPS and XANES spectra at carbon K edge of CuPc are also presented. The comparison between experiment and theory revealed that, in spite of being formally not equivalent, carbon atoms of the benzene rings experience a similar electronic environment. Carbon K-edge absorption spectra were interpreted in terms of different contributions coming from chemically shifted C 1s orbitals of the nonequivalent carbon atoms on the inner ring of the molecule formed by the sequence of CN bonds and on the benzene rings, respectively, and also in terms of different electronic distributions of the excited lowest unoccupied molecular orbital (LUMO) and LUMO+1. In particular, the degenerate LUMO appears to be mostly localized on the inner pyrrole ring.

The role of radial nodes of atomic orbitals for chemical bonding and the periodic table.

Abstract: The role of radial nodes, or of their absence, in valence orbitals for chemical bonding and periodic trends is discussed from a unified viewpoint In particular, we emphasize the special role of the absence of a radial node whenever a shell with angular quantum number l is occupied for the first time (lack of “primogenic repulsion”), as with the 1s, 2p, 3d, and 4f shells Although the consequences of the very compact 2p shell (eg good isovalent hybridization, multiple bonding, high electronegativity, lone-pair repulsion, octet rule) are relatively well known, it seems that some of the aspects of the very compact 3d shell in transition-metal chemistry are less well appreciated, eg, the often weakened and stretched bonds at equilibrium structure, the frequently colored complexes, and the importance of nondynamical electron-correlation effects in bonding © 2006 Wiley Periodicals, Inc J Comput Chem 28: 320–325, 2006

Electronic and vibrational properties of nickel sulfides from first principles

Abstract: We report the results of first-principles calculations (generalized gradient approximation–Perdew Wang 1991) on the electronic and vibrational properties of several nickel sulfides that are observed on Ni-based anodes in solid oxide fuel cells (SOFCs) upon exposure to H2S contaminated fuels: heazlewoodite Ni3S2, millerite NiS, polydymite Ni3S4, and pyrite NiS2. The optimized lattice parameters of these sulfides are within 1% of the values determined from x-ray diffraction. The electronic structure analysis indicates that all Ni–S bonds are strongly covalent. Furthermore, it is found that the nickel d orbitals shift downward in energy, whereas the sulfur p orbitals shift upward with increasing sulfur content; this is consistent with the decrease in conductivity and catalytic activity of sulfur-contaminated Ni-based electrodes (or degradation in SOFC performance). In addition, we systematically analyze the classifications of the vibrational modes at the Γ point from the crystal symmetry and calculate the co...

The Interaction of Dihalogens and Hydrogen Halides with Lewis Bases in the Gas Phase: An Experimental Comparison of the Halogen Bond and the Hydrogen Bond

Abstract: This chapter is concerned exclusively with the experimentally determined properties of halogen-bonded complexes of the type B⋯XY in isolation in the gas phase and their relationship with those of the corresponding hydrogen-bonded complexes B⋯HX. B is one of a series of simple Lewis bases and XY is a homo- or hetero-dihalogen molecule F2, Cl2, Br2, ClF, BrCl or ICl. The method used to determine these properties (angular and radial geometry, binding strength, and the extent of electric charge redistribution on formation of B⋯XY) is first outlined. A comparison of the angular geometries of the pair of halogen-bonded and hydrogen-bonded complexes B⋯ClF and B⋯HCl as B is systematically varied follows. Systematic relationships among the radial geometries of the two series are also summarised. The intermolecular stretching force constants k σ and the extent of electron transfer (both inter- and intramolecular) on formation of B⋯XY, for XY = Cl2, Br2, ClF, BrCl or ICl, are shown to vary systematically as B is varied. A striking similarity noted among the properties of halogen-bonded complexes B⋯XY and their hydrogen-bonded analogues B⋯HX demonstrates that rules for predicting the angular geometries of hydrogen-bonded complexes (and other generalisations) may also be applied to the halogen-bonded series, but with the caveat that while the hydrogen bond shows a propensity to be non-linear when B⋯HX has appropriate symmetry, the halogen bond tends to remain close to linearity. A model for the halogen bond, successfully applied earlier to the hydrogen bond, is proposed.

Gilbert N. Lewis and the chemical bond: The electron pair and the octet rule from 1916 to the present day

Abstract: We describe the development of Lewis's ideas concerning the chemical bond and in particular the concept of the electron pair bond and the octet rule. We show that the concept of the electron pair bond has endured to the present day and is now understood to be a consequence of the Pauli principle. In contrast the octet rule is now regarded as much less important than was originally generally believed, although Lewis himself knew several exceptions and regarded it as less important than what he called the rule of two (the electron pair). The octet rule was more strongly promoted by Langmuir who is also responsible for the term covalent bond. However, many more exceptions to the octet rules than were known to Lewis are now known and the terms hypervalent and hypovalent used to describe such molecules are no longer particularly useful. Today it is realized that bonding electron pairs in many molecules are not as well localized as Lewis believed, nevertheless resonance structures, i.e., plausible alternative Lewis structures, are still often used to describe such molecules. Moreover electrons are not always found in pairs, as for example in linear molecules, which can, however, be satisfactorily described by Linnett's double quartet theory. The electron density distribution in a molecule can now be analyzed using the ELF and other functions of the electron density to show where electron pairs are most probably to be found in a molecule. © 2006 Wiley Periodicals, Inc. J Comput Chem 2007

Size, temperature, and bond nature dependence of elasticity and its derivatives on extensibility, Debye temperature, and heat capacity of nanostructures

Abstract: With the miniaturization of a solid down to nanometer scale, the elasticity, extensibility, Debye temperature, and specific heat capacity of the solid are no longer constant but change with variation of size. These quantities also change with the temperature of the measurement and the nature of the chemical bond involved. The mechanism behind the intriguing tunability and the interdependence of these quantities remain yet a high challenge. A set of analytical solutions is presented herewith showing that the observed trends could be reproduced by taking the fact of bond order deficiency into consideration. Agreement between predictions and observations clarifies that the shortened and strengthened surface bonds dictate intrinsically the observed tunability, yet atoms in the core interior remain as they are in the bulk. The thermally softening of a specimen arises from bond expansion and bond vibration due to the internal energy increases.

Structure and Bonding in β-HMX-Characterization of a Trans-Annular N···N Interaction

Abstract: Chemical bonding in the β-phase of the 1,3,5,7-tetranitro-1,3,5,7-tetraazacyclooctane (HMX) crystal based on the experimental electron density obtained from X-ray diffraction data at 20 K, and solid state theoretical calculations, has been analyzed in terms of the quantum theory of atoms in molecules. Features of the intra- and intermolecular bond critical points and the oxygen atom lone-pair locations are discussed. An unusual N···N bonding interaction across the 8-membered ring has been discovered and characterized. Hydrogen bonding, O···O and O···C intermolecular interactions are reported. Atomic charges and features of the electrostatic potential are discussed.

Effects of doping nitrogen atoms on the structure and electronic properties of zigzag single-walled carbon nanotubes through first-principles calculations

Abstract: Calculations have been made for single-walled zigzag (n, 0) carbon nanotubes containing substitutional nitrogen impurity atoms using ab initio density functional theory. It is found that the formation energies of these nanotubes depend on the tube diameter, as do the electronic properties, and that they show periodic features which result from their different π bonding structures compared to those of perfect zigzag carbon nanotubes. When two nitrogen atoms are doped in the same hexagon per five tube units, the semiconducting tubes exhibit some special electronic structures, in which the impurity level is occupied fully by two excess electrons from doped nitrogen atoms. The electronic structures for the tubes depend on the sites that two nitrogen atoms occupy in the hexagon, by which the impurity states can be near the bottom of the conduction band or can be far apart from the bottom of the conduction band.