Lewis theory electron pairs, importance

However, the protolytic theory cannot explain the distinctly acid or base properties of numerous substances which are not able to either split-off or accept a proton. This stimulated G. N. Lewis (1923) to a different generalization of the notion of acids and bases. According to the Lewis theory a base is a substance which is the donor of a free electron pair, whereas, acid can bond a free electron pair of another particle and thus, it is its acceptor. Neutralization of an acid by base is conditioned by the formation of coordination (donor-acceptor) bond. The Lewis theory is of importance particularly in the chemistry of coordination compounds where all central... 

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 theory as presented so far is clearly incomplete. The topology of the density, while recovering the concepts of atoms, bonds and structure, gives no indication of the localized bonded and non-bonded pairs of electrons of the Lewis model of structure and reactivity, a model secondary in importance only to the atomic model. The Lewis model is concerned with the pairing of electrons, information contained in the electron pair density and not in the density itself. Remarkably enough however, the essential information about the spatial pairing of electrons is contained in the Laplacian of the electron density, the sum of the three second derivatives of the density at each point in space, the quantity V2p(r) [44]. 

The designation of electron-pair donors and acceptors as Lewis bases and Lewis acids is firmly and fittingly ingrained in the language of chemistry. G. N. Lewis laid the foundation for this important theory approximately 80 years ago and Lewis acids have since become increasingly important because of their central role in synthetic organic chemistry. This is clearly illustrated by an ever-increasing number of publications and books. 

In 1902, while explaining the laws of valence to his students at Harvard, Lewis conceived a concrete model for this process, something Abegg had not done. He proposed that atoms were composed of a concentric series of cubes with electrons at each of the resulting eight comers. This cubic atom explained the cycle of eight elements in the Periodic Table and corresponded to the idea that chemical bonds were formed by the transfer of electrons so each atom had a complete set of eight electrons. Lewis did not publish his theory, but fourteen years later it became an important part of his theory on the shared electron-pair bond. 

Nobody will argue the importance of the idea of the electron pair bond, introduced by Lewis [67], in chemistry. Together with the Bohr theory of the electronic structure of the atom [68] and its connection with the periodic system [69], one has the ingredients for a true chemical theory. The octet model introduced by Langmuir [70] soon demonstrated its immense explanative power for organic and inorganic structure alike. 

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. 

In Section 4.3, we learned that the shape of a molecule is an important factor in determining the properties of the substances that it composes. For example, we learned that water would boil away at room temperature if it had a straight shape instead of a bent one. We now develop a simple model called valence shell electron pair repulsion (VSEPR) theory that allows us to predict the shapes of molecules from their Lewis structures. 

The original VSEPR theory (Figure 4.56) is a classical theory that accounts for the shapes of most main group molecules. Lewis structures are simple representations of molecules and ions that emphasize the importance of electron pairs (represented by dots and/or crosses). VSEPR theory proposes that the favoured molecular shape is that which minimizes the repulsion between electron pairs in the valence shell of the central atom. 

It would be wrong to interpret this work as an effort to refute the importance of quantum theory for chemistry. It does the opposite, but questions the methodology that developed from a naive interpretation of three-dimensional wave mechanics to confirm the electron-pair model of Lewis and the molecular structure theory of van t Hoff. Even in terms of the probabilistic interpretation of wave mechanics, a rigid three-dimensionally structured molecule, with its real molecular orbitals, is undefined. A strategy, based on these concepts and which became known as Quantum Chemistry, amounts to a disastrous misreading of quantum theory and has no predictive power beyond its classical basis. 

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