Valence bond theory

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"VBT" redirects here. For other uses, see VBT (disambiguation).
Electronic structure
methods
Valence bond theory
Coulson–Fischer theory
Generalized valence bond
Modern valence bond theory
Molecular orbital theory
Hartree–Fock method
Semi-empirical quantum chemistry methods
Møller–Plesset perturbation theory
Configuration interaction
Coupled cluster
Multi-configurational self-consistent field
Quantum chemistry composite methods
Quantum Monte Carlo
Density functional theory
Time-dependent density functional theory
Thomas–Fermi model
Orbital-free density functional theory
Linearized augmented-plane-wave method
Projector augmented wave method
Electronic band structure
Nearly free electron model
Tight binding
Muffin-tin approximation
k·p perturbation theory
Empty lattice approximation
GW approximation
Korringa–Kohn–Rostoker method

In chemistry, valence bond (VB) theory is one of the two basic theories, along with molecular orbital (MO) theory, that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the atomic orbitals of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, molecular orbital theory has orbitals that cover the whole molecule.[1]

History

In 1916,

ionic bonding. Both Lewis and Kossel structured their bonding models on that of Abegg's rule
(1904).

Although there is no mathematical formula either in chemistry or quantum mechanics for the arrangement of electrons in the atom, the hydrogen atom can be described by the

orbital hybridization (1930). According to Charles Coulson, author of the noted 1952 book Valence, this period marks the start of "modern valence bond theory", as contrasted with older valence bond theories, which are essentially electronic theories of valence
couched in pre-wave-mechanical terms.

Linus Pauling published in 1931 his landmark paper on valence bond theory: "On the Nature of the Chemical Bond". Building on this article, Pauling's 1939 textbook: On the Nature of the Chemical Bond would become what some have called the bible of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to adequately address the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960s and 1970s as molecular orbital theory grew in usefulness as it was implemented in large

digital computer
programs. Since the 1980s, the more difficult problems, of implementing valence bond theory into computer programs, have been solved largely, and valence bond theory has seen a resurgence.

Theory

According to this theory a covalent bond is formed between two atoms by the overlap of half filled valence atomic orbitals of each atom containing one unpaired electron. A valence bond structure is similar to a

electrons should be in the bond region. Valence bond theory views bonds as weakly coupled orbitals (small overlap). Valence bond theory is typically easier to employ in ground state molecules. The core orbitals and electrons
remain essentially unchanged during the formation of bonds.

σ bond between two atoms: localization of electron density
Two p-orbitals forming a π-bond.

The overlapping atomic orbitals can differ. The two types of overlapping orbitals are sigma and pi.

hybridization
.

Comparison with MO theory

Valence bond theory complements

dihydrogen
is an equal mixture of the covalent and ionic valence bond structures and so predicts incorrectly that the molecule would dissociate into an equal mixture of hydrogen atoms and hydrogen positive and negative ions.

Computational approaches

Main article: Modern valence bond theory

Modern valence bond theory replaces the overlapping atomic orbitals by overlapping valence bond orbitals that are expanded over a large number of

Hartree–Fock reference wavefunction. The most recent text is by Shaik and Hiberty.[10]

Applications

An important aspect of the valence bond theory is the condition of maximum overlap, which leads to the formation of the strongest possible bonds. This theory is used to explain the covalent bond formation in many molecules.

For example, in the case of the F2 molecule, the F−F bond is formed by the overlap of pz orbitals of the two F atoms, each containing an unpaired electron. Since the nature of the overlapping orbitals are different in H2 and F2 molecules, the bond strength and bond lengths differ between H2 and F2 molecules.

In an HF molecule the covalent bond is formed by the overlap of the 1s orbital of H and the 2pz orbital of F, each containing an unpaired electron. Mutual sharing of electrons between H and F results in a covalent bond in HF.

Using modern classical valence bond theory, Patil and Bhanage have shown that the cation-anion interface of protic ionic liquids possesses charge shift bond character.[11]

See also

References

  1. ISBN 0-471-90759-6
    .
  2. ISSN 0002-7863
    .
  3. ^ University College Cork, University City Tübingen, and (Pauling, 1960, p. 5).
  4. ^ Walther Kossel, “Uber Molkulbildung als Frage der Atombau”, Ann. Phys., 1916, 49:229–362.
  5. ^ Walter Heitler – Key participants in the development of Linus Pauling's The Nature of the Chemical Bond.
  6. S2CID 24349360
    .
  7. S2CID 4261220
    .
  8. S2CID 45218186
    .
  9. S2CID 4268597
    .
  10. ISBN 978-0-470-03735-5
    .
  11. PMID 27229870
    . Retrieved 25 June 2022.
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