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London atomic orbitals chemical

Lee AM, Handy NC, Colwell SM (1995) The density functional calculation of nuclear shielding constants using London atomic orbitals. J ChemPhys 103 10095-10109 Lippmaa,E, Magi M, Samoson A, Engelhard G, Grimmer,A-R (1980) Structural studies of silicates by solid-state high-resolution Si NMR. J Am Chem Soc 102 4889-4893 Lipscomb WN (1966) The chemical shift and other second-order magnetic and electric properties of small molecules. In Adv Mag Reson, Vol. 2, p 137-176... [Pg.456]

Thorvaldsen, A. J., Ruud, K., Rizzo, A., 8c Coriani, S. (2008). Analytical calculations of frequency-dependent hypermagnetizabilities and Cotton-Mouton constants using London atomic orbitals. Journal of Chemical Physics, 129, 164110. [Pg.440]

N. N. Greenwood, Principles of Atomic Orbitals, revised SI edition. Monograph for Teachers, No. 8, Chemical Society, London, 1980, 48 pp. [Pg.21]

The first calculations on a two-electron bond was undertaken by Heitler and London for the H2 molecule and led to what is known as the valence bond approach. While the valence bond approach gained general acceptance in the chemical community, Robert S. Mulliken and others developed the molecular orbital approach for solving the electronic structure problem for molecules. The molecular orbital approach for molecules is the analogue of the atomic orbital approach for atoms. Each electron is subject to the electric field created by the nuclei plus that of the other electrons. Thus, one was led to a Hartree-Fock approach for molecules just as one had been for atoms. The molecular orbitals were written as linear combinations of atomic orbitals (i.e. hydrogen atom type atomic orbitals). The integrals that needed to be calculated presented great difficulty and the computations needed were... [Pg.51]

Fig. 18. (a) Orbital overlap scheme for three-center P— N P bonds. (Shaded atomic orbitals are combined in molecular orbitals.) (Reproduced from Craig and Mitchell (135a) by permission of The Chemical Society, London.) (b) Three-center P—N—P islands in N3P,R6. [Pg.96]

Hall G G 1951 The Molecular Orbital Theory of Chemical Valency VIII A Method for Calculating lomsation Potentials. Proceedings of the Royal Society (London) A205 541-552 Hehre W J, R F Stewart and J A Pople 1%9 Self-Consistent Molecular-Orbital Methods I Use of Gaussian Expansions of Slater-Type Atomic Orbitals. Journal of Chemical Physics 51 2657-2664 Hehre W J, L Radom, P v R Schleyer and J A Pople 1986 Ah initio Molecular Orbital Theory New York, John Wiley Sons. [Pg.106]

The most intuitively familiar model of molecular electronic structure which gives a satisfactory qualitative and quantitative picture of both stable bonds and bondbreaking is the famous Heitler-London non-polar covalent model. A chemical bond between a pair of (hybrid atomic) orbitals Aj and A2 is described by the two-electron wavefunction... [Pg.289]

The first quantitative theory of chemical bonding was developed for the hydrogen molecule by Heitler and London in 1927, and was based on the Lewis theory of valence in which two atoms shared electrons in such a way that each achieved a noble gas structure. The theory was later extended to other, more complex molecules, and became known as valence bond theory. In this approach, the overlap of atomic orbitals on neighbouring atoms is considered to lead to the formation of localized bonds, each of which can accommodate two electrons with paired spins. The theory has been responsible for introducing such important concepts as hybridization and resonance into the theory of the chemical bond, but applications of the theory have been limited by difficulties in generating computer programs that can deal efficiently with anything other than the simplest of molecules. [Pg.137]

In 1927, Walter Heitler and Fritz Wolfgang London clarified the origin of the covalent chemical bond (Heitler and London 1927), the concept crucial for chemistry. In the paper the authors demonstrated, in numerical calculations, that the nature of the covalent chemical bond in H2 is of quantum character, because the (semiquantitatively) correct description of H2 emerged only after inclusion the exchange of electrons 1 and 2 between the nuclei in the formula a l)b 2) (a, b are the Is atomic orbitals centered on nucleus a and nucleus b, respectively) resulting in the wave function a(l)b 2)+a 2)b(l). Thus, taking into account also the contribution of Hund (1927a, b, c), 1927 is therefore the year of birth of computational chemistry. [Pg.4]

It will be recalled that the first quantum-mechanical explanation of the chemical bond is usually attributed to Heitler and London (1927). Their discussion of the hydrogen molecule, however, was not based on IPM concepts, with each electron assigned to a molecular orbital, but rather on an independent-atom approach in which the electrons were assigned to atomic orbitals. The Heitler-London approach was the basis for valence bond (VB) theory, which was important in the early days of quantum chemistry but later fell into disuse. There has been some revival of interest in VB theory now that more powerful computing facilities are available a full discussion of this method is deferred until Chapter 7, although a preliminary account will appear in Chapter 3. [Pg.20]

In the late 1920s, it was shown that the chemical bond existing between two identical hydrogen atoms in H2 can be described mathematically by taking a linear combination of the Is orbitals [Pg.176]

Additional information on the history of atomic theory can be found in J. R. Partington, A Short History of Chemistry, 3rd ed., Macmillan, London, 1957, reprinted by Harper Row, New York, 1960, and in the Journal of Chemical Education. A more thorough treatment of the electronic structure of atoms is in M. Gerloch, Orbitals, Terms, and States, John Wiley Sons, New York, 1986. [Pg.47]

The basic idea of the Heitler-London model for the hydrogen molecule can be extended to chemical bonds between any two atoms. The orbital function (10.8) must be associated with the singlet spin function cro,o(l > 2) in order that the overall wavefunction be antisymmetric [cf. Eq (8.14)]. This is a quantum-mechanical realization of the concept of an electron-pair bond, first proposed by G. N. Lewis in 1916. It is also now explained why the electron spins must be paired, i.e., antiparallel. It is also permissible to combine an antisymmetric orbital function with a triplet spin function, but this will, in most cases, give a repulsive state, such as the one shown in red in Fig. 10.2. [Pg.77]


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