Chapter 2: Molecular Structure and Bonding Bonding Theories (original) (raw)
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Molecular orbital theory is a conceptual extension of the orbital model, which was so successfully applied to atomic structure. As was once playfully remarked, "a molecule is nothing more than an atom with more nuclei." This may be overly simplistic, but we do attempt, as far as possible, to exploit analogies with atomic structure. Our understanding of atomic orbitals began with the exact solutions of a prototype problem – the hydrogen atom. We will begin our study of homonuclear diatomic molecules beginning with another exactly solvable prototype, the hydrogen molecule-ion \(H_{2}^{+}\).
New bonding concept for Hypervalent molecules, including electron poor and electron odd compounds
For the question about the real bond's nature of so told Hypervalent molecules, the new concept (scheme) for electron arrangement in an atoms and molecules in next article is presented. This concept is not something new, as new theory, but only new come across on chemical formulas and data from many scientific articles, literature and internet, altogether with current actual scientific laws and theories. The base for this concept is to use the real energy sublevel order instead quantum mechanical sublevel order in the electron configuration of the elements.. This concept than will became powerful tool for solving the structure and properties of many regular and irregular molecules or compounds. Same as Lewis octet rule, which rule is mostly establish on conclusion that in most stable molecules atoms have eight electrons in their valence shell, this new concept is also acquired by observation (conclusion) that data and chemical formulas for many compounds can be presented in different way that it is currently done. All of this leads to the new presentation of well-known facts of electron arrangement in sublevels. This new electron arrangement is accomplished with the new pattern of the electron box diagrams, named as arrow-dot orbital box diagrams. In addition, here, several new important principles are established. These new diagrams and principles are very good tool for researches and especially for chemist because with them, chemical bonding and their diagrams, for many unusual and unsolved molecules, precise and with more certainty can be explained and created.
Valence Bond and Molecular Orbital: Two Powerful Theories that Nicely Complement One Another
Journal of Chemical Education, 2021
Introductory chemistry textbooks often present valence bond (VB) theory as useful, but incorrect and inferior to molecular orbital (MO) theory, citing the electronic structure of O 2 and electron delocalization as evidence. Even texts that initially present the two theories on equal footing use language that biases students toward the MO approach. However, these "failures" of VB are really just misconceptions and/or misapplications of the theory. At their theoretical limits, both VB and MO are equivalent; they simply approach that limit from different sides. Certain concepts may be easier to grasp with one theory or the other so that having a commanding knowledge of both is extremely beneficial. However, presenting one theory as superior to the other suppresses the ability to look at a problem from both sides and is therefore detrimental to students and the whole of chemistry. It is time for VB and MO to be taught on equal footing like the complementary theories they are.
Bond electron pair: Its relevance and analysis from the quantum chemistry point of view
Journal of Computational Chemistry, 2007
This paper first comments on the surprisingly poor status that Quantum Chemistry has offered to the fantastic intuition of Lewis concerning the distribution of the electrons in the molecule. Then, it advocates in favor of a hierarchical description of the molecular wave-function, distinguishing the physics taking place in the valence space (in the bond and between the bonds), and the dynamical correlation effects. It is argued that the clearest pictures of the valence electronic population combine two localized views, namely the bond (and lone pair) Molecular Orbitals and the Valence Bond decomposition of the wave-function, preferably in the orthogonal version directly accessible from the complete active space self consistent field method. Such a reading of the wave function enables one to understand the work of the nondynamical correlation as an enhancement of the weight of the low-energy VB components, i.e. as a better compromise between the electronic delocalization and the energetic preferences of the atoms. It is suggested that regarding the bond building, the leading dynamical correlation effect may be the dynamical polarization phenomenon. It is shown that most correlation effects do not destroy the bond electron pairs and remain compatible with Lewis' vision. A certain number of free epistemological considerations have been introduced in the development of the argument. q
On the electron-pair nature of the hydrogen bond in the framework of the atoms in molecules theory
2003
Delocalization indices, as defined in the atoms in molecules theory, have been calculated between hydrogen-bonded atoms in 20 molecular complexes that are formed between several H-donor and acceptor molecules. In general, the delocalization index associated to an intermolecular hydrogen bond depends on the interaction energy of the complex, but also on the nature of the H-donor and acceptor atoms.
Appendix 10: The Origin of Chemical Bonding in H2+
2009
In Chapter 7 of the main text we outline the various contributions to the stabilization energy of the simplest of all molecules: the H 2 + cation. The main goal there is to introduce the concepts of MO theory to use in more complex molecules. In this appendix, some additional notes are given on the role of the various components of the kinetic and potential energy terms that contribute to the chemical bond energy of the H 2 + cation. We will begin by following the general idea that a linear combination of AOs can be used to form MOs by taking the 1s functions of the H atom derived in Appendix 9 as our basis. Half way through this discussion we will come to the conclusion that chemical bonds decrease the electron's kinetic energy but increase its potential energy. However, we will find that the energies calculated violate an important theorem in the expected balance of average kinetic and potential energy for the molecular ion. This demonstrates that the linear combination of AOs using the radial decay factors for atoms cannot give a complete picture of chemical bond formation. The radial profiles of the orbitals also have to be allowed to adapt to account for the changing environment the electron experiences on moving from the AOs to MOs. Once this is done, the potential energy is decreased and the kinetic energy increased due to the contraction of the orbitals around the nuclei. The flexibility of the radial shape of the AOs used in the description of MOs is an important one in theoretical chemistry and will be discussed further in Appendix 11.
On The Nature of the Chemical Bond in Valence Bond Theory
The Journal of Chemical Physics
This perspective outlines a panoramic description of the nature of the chemical bond according to valence bond theory. It describes single bonds, and charge-shift bonds (CSBs) in which the entire/most of the bond energy arises from the resonance between the covalent and ionic structures of the bond. Many CSBs are homonuclear bonds. Hypervalent molecules are CSBs. Then we describe multiply bonded molecules with emphasis on C2 and 3O2. The perspective outlines an effective methodology of peeling the electronic structure to the necessary minimum: a structure with a quadruple bond, and two minor structures with double bonds, which stabilize the quadruple bond by resonance. 3O2 is chosen because it is a persistent diradical. The persistence of 3O2 is due to the large CSB resonance interaction of the π-3-electron bonds. Subsequently, we describe the roles of π vs. σ in the geometric preferences in unsaturated molecules, and their Si-based analogs. Then, the perspective discusses bonding i...
International Journal of Quantum Chemistry, 2018
The derivative of molecular orbitals (MO) energies with respect to a bond length (dynamic orbital force, DOF) is used to estimate the bonding/antibonding character of valence MOs along this bond, with a focus on lone pair MOs, in a series of small molecules: AH (A = F, Cl, Br), AH2 (A = O, S, Se) AX3 (A = N, P, As; X = H, F) and H2CO. The HOMO DOF agrees with the calculated variation of bond length and force constant in the corresponding ground state cation, and of bond length variation by protonation. These results also agree with available experimental data. It is worthy to note that the p-type HOMOs in AH and AH2 are found bonding. The lone pair MO is bonding in NH3, while it is antibonding in PH3, AsH3, and AF3.
Electron pairing and chemical bonds. On the accuracy of the electron pair model of chemical bond
Journal of Molecular Structure: THEOCHEM, 1997
The accuracy of the Lewis electron pair model of the chemical bond is reinvestigated. It is shown that, contrary to previous pessimistic findings, the accuracy of this model is high enough to represent a good basis for the understanding and interpretation of molecular structure, not only for molecules well described by a classical structural formula with localised two-centre hvoelectron (2c-2e) bonds, but also for molecules with more complex bonding patterns such as three-or multicentre bonding. Keywords: Electron pairing; Lewis model of chemical bond; Pair population analysis 0166-12&X0/97/$17.00 Copyright 0 1997 Elsevier Science B.V. All rights reserved Plf SO166-1280(96)04728-S