Atoms, bonded pairs

Consider all valence electron pairs of the central atom-bonding pairs,... 

Determine the total number of electron groups around the central atom (bonding pairs, lone pairs and, where applicable, account for the charge on the ion). Remember that a double bond or a triple bond is counted as one electron group. 

Count the number of electron pairs around the central atom (bonding pairs and lone pairs). Treat double and triple bonds as though they were single bonds. Refer to Table 10.1 to predict the overall arrangement of the electron pairs. 

Bond enei Amount of energy required to break one mole of bonds between a specified pair of atoms Bonding pairs Pairs of electrons shared between two atoms, frequentiy represented by lines (instead of dots) for convenience in Lewis structures... 

The concept of resonance is an adaptation of the Lewis model that helps account for the complexity of actual molecules. In the Lewis model, electrons are localized either on one atom (lone pair) or between atoms (bonding pair). However, in nature, the electrons in molecules are often delocalized over several atoms or bonds. The delocalization of electrons lowers their energy it stabilizes them (for reasons that are beyond the scope of this book). Resonance depicts two or more structures with the electrons in different places in an attempt to more accurately reflect the delocalization of electrons. In the real hybrid structure, an average between the resonance structures, the electrons are more spread out (or delocalized) than in any of the resonance structures. The resulting stabilization of the electrons (that is, the lowering of their potential energy due to delocalization) is sometimes called resonance stabilization. Resonance stabilization makes an important contribution to the stability of many molecules. 

The iHoblem of graph symmetry investigates the equivalence relationships between the elements of the molecular graphs (atoms, bonds, pairs of atoms, etc.). The geometrical information is neglected and only bonding relationships are considered. 

A simple example would be in a study of a diatomic molecule that in a Hartree-Fock calculation has a bonded cr orbital as the highest occupied MO (HOMO) and a a lowest unoccupied MO (LUMO). A CASSCF calculation would then use the two a electrons and set up four CSFs with single and double excitations from the HOMO into the a orbital. This allows the bond dissociation to be described correctly, with different amounts of the neutral atoms, ion pair, and bonded pair controlled by the Cl coefficients, with the optimal shapes of the orbitals also being found. For more complicated systems... 

Methane, CH4, for example, has a central carbon atom bonded to four hydrogen atoms and the shape is a regular tetrahedron with a H—C—H bond angle of 109°28, exactly that calculated. Electrons in a lone pair , a pair of electrons not used in bonding, occupy a larger fraction of space adjacent to their parent atom since they are under the influence of one nucleus, unlike bonding pairs of electrons which are under the influence of two nuclei. Thus, whenever a lone pair is present some distortion of the essential shape occurs. 

In some force fields the interaction sites are not all situated on the atomic nuclei. For example, in the MM2, MM3 and MM4 programs, the van der Waals centres of hydrogen atoms bonded to carbon are placed not at the nuclei but are approximately 10% along the bond towards the attached atom. The rationale for this is that the electron distribution about small atoms such as oxygen, fluorine and particularly hydrogen is distinctly non-spherical. The single electron from the hydrogen is involved in the bond to the adjacent atom and there are no other electrons that can contribute to the van der Waals interactions. Some force fields also require lone pairs to be defined on particular atoms these have their own van der Waals and electrostatic parameters. 

The tetrahedral geometry of methane is often explained with the valence shell electron pair repulsion (VSEPR) model The VSEPR model rests on the idea that an electron pair either a bonded pair or an unshared pair associated with a particular atom will be as far away from the atom s other electron pairs as possible Thus a tetrahedral geomehy permits the four bonds of methane to be maximally separated and is charac terized by H—C—H angles of 109 5° a value referred to as the tetrahedral angle... 

The reactivity of the individual O—P insecticides is determined by the magnitude of the electrophilic character of the phosphoms atom, the strength of the bond P—X, and the steric effects of the substituents. The electrophilic nature of the central P atom is determined by the relative positions of the shared electron pairs, between atoms bonded to phosphoms, and is a function of the relative electronegativities of the two atoms in each bond (P, 2.1 O, 3.5 S, 2.5 N, 3.0 and C, 2.5). Therefore, it is clear that in phosphate esters (P=0) the phosphoms is much more electrophilic and these are more reactive than phosphorothioate esters (P=S). The latter generally are so stable as to be relatively unreactive with AChE. They owe their biological activity to m vivo oxidation by a microsomal oxidase, a reaction that takes place in insect gut and fat body tissues and in the mammalian Hver. A typical example is the oxidation of parathion (61) to paraoxon [311-45-5] (110). 

This initial assignment is, of course, not at equilibrium. In particular, the expected velocity correlation between neighboring atoms is not guaranteed, and most likely it is nonexistent (i.e., in general, neighboring atoms, such as bonded pairs, are expected to... 

At room temperature VO2 has a rutile-like strucmre (p. 961) distorted by the presence of pairs of vanadium atoms bonded together. Above 70°C, however, an undistorted rutile... 

Off-diagonal elements for bonded pairs of atoms depend only on the types of atom involved. 

Structural Isomers. These contain different combinations of bonded pairs of atoms. They may be divided into three types chain, position, and functional isomers,... 

One further point inductive effects and resonance effects don t necessarily act in the same direction. Halogen, hydroxyl, alkoxyl, and amino substituents, for instance, have electron -withdrawing inductive effects because of the electronegativity of the -X, -O, or —N atom bonded to the aromatic ring but have resonance effects because of the lone-pair electrons on those same —X, -O, or —N atoms. When the two effects act in opposite directions, the stronger of the two dominates. 

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