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Orbital Overlap and Chemical Bonding
Author(s) -
Krapp Andreas,
Bickelhaupt F. Matthias,
Frenking Gernot
Publication year - 2006
Publication title -
chemistry – a european journal
Language(s) - English
Resource type - Journals
SCImago Journal Rank - 1.687
H-Index - 242
eISSN - 1521-3765
pISSN - 0947-6539
DOI - 10.1002/chem.200600564
Subject(s) - pauli exclusion principle , valence bond theory , atomic orbital , molecular orbital theory , molecular orbital , chemistry , molecular orbital diagram , orbital hybridisation , non bonding orbital , atomic physics , chemical bond , covalent bond , pi bond , orbital overlap , physics , molecular physics , electron , bond length , molecule , bond order , quantum mechanics , organic chemistry
The chemical bonds in the diatomic molecules Li 2 –F 2 and Na 2 –Cl 2 at different bond lengths have been analyzed by the energy decomposition analysis (EDA) method using DFT calculations at the BP86/TZ2P level. The interatomic interactions are discussed in terms of quasiclassical electrostatic interactions Δ E elstat , Pauli repulsion Δ E Pauli and attractive orbital interactions Δ E orb . The energy terms are compared with the orbital overlaps at different interatomic distances. The quasiclassical electrostatic interactions between two electrons occupying 1s, 2s, 2p(σ), and 2p(π) orbitals have been calculated and the results are analyzed and discussed. It is shown that the equilibrium distances of the covalent bonds are not determined by the maximum overlap of the σ valence orbitals, which nearly always has its largest value at clearly shorter distances than the equilibrium bond length. The crucial interaction that prevents shorter bonds is not the loss of attractive interactions, but a sharp increase in the Pauli repulsion between electrons in valence orbitals. The attractive interactions of Δ E orb and the repulsive interactions of Δ E Pauli are both determined by the orbital overlap. The net effect of the two terms depends on the occupation of the valence orbitals, but the onset of attractive orbital interactions occurs at longer distances than Pauli repulsion, because overlap of occupied orbitals with vacant orbitals starts earlier than overlap between occupied orbitals. The contribution of Δ E elstat in most nonpolar covalent bonds is strongly attractive. This comes from the deviation of quasiclassical electron–electron repulsion and nuclear–electron attraction from Coulomb's law for point charges. The actual strength of Δ E elstat depends on the size and shape of the occupied valence orbitals. The attractive electrostatic contributions in the diatomic molecules Li 2 –F 2 come from the s and p(σ) electrons, while the p(π) electrons do not compensate for nuclear–nuclear repulsion. It is the interplay of the three terms Δ E orb , Δ E Pauli , and Δ E elstat that determines the bond energies and equilibrium distances of covalently bonded molecules. Molecules like N 2 and O 2 , which are usually considered as covalently bonded, would not be bonded without the quasiclassical attraction Δ E elstat .