Molecules orbital theory

The crystal structure of 864(118207)2 90, 91) has shown 864 to be square planar with an 8e-8e bond distance of 2.283(4) A, significantly less than that of 2.34(2) A found in the 8eg molecule (92), indicating some degree of multiple bonding. 8uch a result is consistent with a valence bond description of the molecule involving four structures of type VII. Alternatively the structure can be understood in terms of molecule orbital theory. The circle in structure VIII denotes a closed-shell (aromatic ) six-w-electron system. Of the four tt molecular orbitals,... 

Periodic system of elements Structure of molecules (orbital theory)... 

J. N. Murrell, A. J. Harget, Semi-empirical self-consistent-field molecular orbital theory of molecules John Wiley Sons, New York (1972). 

Orbital Theories of Molecules and Solids N. H. March, Ed., Clarendon, Oxford (1974). 

Valence bond and molecular orbital theory both incorporate the wave description of an atom s electrons into this picture of H2 but m somewhat different ways Both assume that electron waves behave like more familiar waves such as sound and light waves One important property of waves is called interference m physics Constructive interference occurs when two waves combine so as to reinforce each other (m phase) destructive interference occurs when they oppose each other (out of phase) (Figure 2 2) Recall from Section 1 1 that electron waves m atoms are characterized by their wave function which is the same as an orbital For an electron m the most stable state of a hydrogen atom for example this state is defined by the Is wave function and is often called the Is orbital The valence bond model bases the connection between two atoms on the overlap between half filled orbifals of fhe fwo afoms The molecular orbital model assembles a sef of molecular orbifals by combining fhe afomic orbifals of all of fhe atoms m fhe molecule... 

Section 2 4 In molecular orbital theory the molecular orbitals (MOs) are approxi mated by combining the atomic orbitals (AOs) of all of the atoms m a molecule The number of MOs must equal the number of AOs that are combined... 

Molecular ion (Section 13 22) In mass spectrometry the species formed by loss of an electron from a molecule Molecular orbital theory (Section 2 4) Theory of chemical bonding in which electrons are assumed to occupy orbitals in molecules much as they occupy orbitals in atoms The molecular orbitals are descnbed as combinations of the or bitals of all of the atoms that make up the molecule Molecularity (Section 4 8) The number of species that react to gether in the same elementary step of a reaction mechanism... 

Frontier orbital theory also provides the basic framework for analysis of the effect that the symmetiy of orbitals has upon reactivity. One of the basic tenets of MO theory is that the symmetries of two orbitals must match to permit a strong interaction between them. This symmetry requirement, when used in the context of frontier orbital theory, can be a very powerful tool for predicting reactivity. As an example, let us examine the approach of an allyl cation and an ethylene molecule and ask whether the following reaction is likely to occur. 

For a molecule as simple as Fl2, it is hard to see much difference between the valence bond and molecular orbital methods. The most important differences appear- in molecules with more than two atoms. In those cases, the valence bond method continues to view a molecule as a collection of bonds between connected atoms. The molecular- orbital method, however, leads to a picture in which the sane electron can be associated with many, or even all, of the atoms in a molecule. We ll have more to say about the similarities and differences in valence bond and molecular- orbital theory as we continue to develop their principles, beginning with the simplest alkanes methane, ethane, and propane. 

Benzene is described by molecular orbital theory as a planar, cyclic, conjugated molecule with six it electrons. According to the Htickel rule, a molecule must have 4n + 2 77 electrons, where n - 0, 1, 2, 3, and so on, to be aromatic. Pianar, cyclic, conjugated molecules with other numbers of tt electrons are antiaromatic. 

This discrepancy between experiment and theory (and many others) can be explained in terms of an alternative model of covalent bonding, the molecular orbital (MO) approach. Molecular orbital theory treats bonds in terms of orbitals characteristic of the molecule as a whole. To apply this approach, we carry out three basic operations. 

To illustrate molecular orbital theory, we apply it to the diatomic molecules of the elements in the first two periods of the periodic table. 

Among the diatomic molecules of the second period elements are three familiar ones, N2,02, and F2. The molecules Li2, B2, and C2 are less common but have been observed and studied in the gas phase. In contrast, the molecules Be2 and Ne2 are either highly unstable or nonexistent. Let us see what molecular orbital theory predicts about the structure and stability of these molecules. We start by considering how the atomic orbitals containing the valence electrons (2s and 2p) are used to form molecular orbitals. 

Hurley, A. C., Proc. Roy. Soc. [London) A216, 424, The molecular orbital theory of chemical valency. XIII. Orbital wave functions for excited states of a homonuclear diatomic molecule."... 

Thus the orbitals and r electrons lie in the outermost part of the valence shell of ethane. They should play a critical role in determining the chemical properties of the molecule. Some theories have ascribed the barrier to internal rotation to these orbitals. It should be noted that the existence of r electrons in ethane is not a novelty, and was first pointed out by Mullikcn in 1935. 

Lewis s theory of the chemical bond was brilliant, but it was little more than guesswork inspired by insight. Lewis had no way of knowing why an electron pair was so important for the formation of covalent bonds. Valence-bond theory explained the importance of the electron pair in terms of spin-pairing but it could not explain the properties of some molecules. Molecular orbital theory, which is also based on quantum mechanics and was introduced in the late 1920s by Mul-liken and Hund, has proved to be the most successful theory of the chemical bond it overcomes all the deficiencies of Lewis s theory and is easier to use in calculations than valence-bond theory. 

The development of molecular orbital theory (MO theory) in the late 1920s overcame these difficulties. It explains why the electron pair is so important for bond formation and predicts that oxygen is paramagnetic. It accommodates electron-deficient compounds such as the boranes just as naturally as it deals with methane and water. Furthermore, molecular orbital theory can be extended to account for the structures and properties of metals and semiconductors. It can also be used to account for the electronic spectra of molecules, which arise when an electron makes a transition from an occupied molecular orbital to a vacant molecular orbital. 

In molecular orbital theory, electrons occupy orbitals called molecular orbitals that spread throughout the entire molecule. In other words, whereas in the Lewis and valence-bond models of molecular structure the electrons are localized on atoms or between pairs of atoms, in molecular orbital theory all valence electrons are delocalized over the whole molecule, not confined to individual bonds. 

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