Molecular orbitals repulsion approach

The molecular orbital (MO) approach developed more slowly but was ultimately much more successful than the other approaches. Dewar (1950) was able to predict absorption maxima for various types of cyanine dyes in excellent agreement with experiment. A major advance was made in 1953 when a self-consistent molecular orbital method specifically taking into account antisymmetrisation and electron repulsion effects was developed by Pariser, Parr and Pople. The PPP-MO method established itself over the next two decades as the most useful and versatile technique for colour prediction, especially when the microelectronics revolution provided facilities for overcoming the complexity of the necessary calculations. 

A cyclic transition state model, that differs from the Zimmerman-Traxler and the related cyclic models inasmuch as it does not incorporate the metal in a chelate, has been proposed by Mulzer and coworkers [78] It has been developed as a rationale for the observation that, in the aldol addition of the dianion of phenylacetic acid 152, the high ti-selectivity is reached with naked enolate anions (e.g., with the additive 18-crown-6). Thus, it was postulated that the approach of the enolate to the aldehyde is dominated by an interaction of the enolate HOMO and the n orbital of the aldehyde that functions as the LUMO (Scheme 4.31), the phenyl substituents of the enolate (phenyl) and the residue R of the aldehyde being oriented in anti position at the forming carbon bond, so that the steric repulsion in the transition state 153 is minimized. Mulzer s frontier molecular orbital-inspired approach reminds of a 1,3-dipolar cycloaddition. However, the corresponding cycloadduct 154 does not form, because of the weakness of the oxygen-oxygen bond. Instead, the doubly metallated aldol adduct 155 results. Anh and coworkers also emphasized the frontier orbital interactions as being essential for the stereochemical outcome of the aldol reaction [79]. 

CH3I should approach the enolate from the direction that simultaneously allows its optimum overlap with the electron-donor orbital on the enolate (this is the highest-occupied molecular orbital or HOMO), and minimizes its steric repulsion with the enolate. Examine the HOMO of enolate A. Is it more heavily concentrated on the same side of the six-membered ring as the bridgehead methyl group, on the opposite side, or is it equally concentrated on the two sides A map of the HOMO on the electron density surface (a HOMO map ) provides a clearer indication, as this also provides a measure of steric requirements. Identify the direction of attack that maximizes orbital overlap and minimizes steric repulsion, and predict the major product of each reaction. Do your predictions agree with the thermodynamic preferences Repeat your analysis for enolate B, leading to product B1 nd product B2. 

For planar unsaturated and aromatic molecules, many MO calculations have been made by treating the a and n electrons separately. It is assumed that the o orbitals can be treated as localized bonds and the calculations involve only the tt electrons. The first such calculations were made by Hiickel such calculations are often called Hiickel molecular orbital (HMO) calculations Because electron-electron repulsions are either neglected or averaged out in the HMO method, another approach, the self-consistent field (SCF), or Hartree-Fock (HF), method, was devised. Although these methods give many useful results for planar unsaturated and aromatic molecules, they are often unsuccessful for other molecules it would obviously be better if all electrons, both a and it, could be included in the calculations. The development of modem computers has now made this possible. Many such calculations have been made" using a number of methods, among them an extension of the Hiickel method (EHMO) and the application of the SCF method to all valence electrons. ... 

Other approximate, more empirical methods are the extended Huckel 31> and hybrid-based Hiickel 32. 3> approaches. In these methods the electron repulsion is not taken into account explicitly. These are extensions of the early Huckel molecular orbitals 4> which have successfully been used in the n electron system of planar molecules. On account of the simplest feature of calculation, the Hiickel method has made possible the first quantum mechanical interpretation of the classical electronic theory of organic chemistry and has given a reasonable explanation for the chemical reactivity of sizable conjugated molecules. 

Like atomic orbitals (AOs), molecular orbitals (MOs) are conveniently described by quantum mechanics theory. Nevertheless, the approach is more complex, because the interaction involves not simply one proton and one electron, as in the case of AOs, but several protons and electrons. For instance, in the simple case of two hydrogen atoms combined in a diatomic molecule, the bulk coulombic energy generated by the various interactions is given by four attractive effects (proton-electron) and two repulsive effects (proton-proton and electron-electron cf figure 1.20) ... 

Exact solutions such as those given above have not yet been obtained for the usual many-electron molecules encountered by chemists. The approximate method which retains tile idea of orbitals for individual electrons is called molecular-orbital theory (M. O. theory). Its approach to the problem is similar to that used to describe atomic orbitals in the many-electron atom. Electrons are assumed to occupy the lowest energy orbitals with a maximum population of two electrons per orbital (to satisfy the Pauli exclusion principle). Furthermore, just as in the case of atoms, electron-electron repulsion is considered to cause degenerate (of equal energy) orbitals to be singly occupied before pairing occurs. 

We now have four molecular orbitals, o, cr2, crx and cry, one lowered in energy and one raised relative to the energy of the orbitals of the pair of hydrogen molecules. If we have four electrons in the system, the net result is repulsion. Thus two H2 molecules do not combine to form an H4 molecule. This is true whatever geometry we use in the combination. It shows us why molecules exist—when two molecules approach each other, the interaction of their molecular orbitals usually leads to repulsion. 

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