Bonds shared-electron-pair

Connect bonded atoms by a shared electron pair bond () represented by a dash (—)... 

Count the number of electrons in shared electron pair bonds (twice the number of bonds) and sub tract this from the total number of electrons to give the number of electrons to be added to com plete the structure... 

Orbital hybridization descriptions because they too are based on the shared electron pair bond enhance the information content of Lewis formulas by distinguishing... 

Nucleophilic addition (Section 17 6) The charactenstic reac tion of an aldehyde or a ketone An atom possessing an un shared electron pair bonds to the carbon of the C=0 group and some other species (normally hydrogen) bonds to the oxygen... 

Polar covalent bond (Section 1.5) A shared electron pair bond in which the electrons are drawn more closely to one of the bonded atoms than the other. 

An equation has been formulated to express the change in covalent radius (metallic radius) of an atom with change in bond number (or in coordination number, if the valence remains constant), the stabilizing (bond-shortening) effect of the resonance of shared-electron-pair bonds among alternative positions being also taken into consideration. This equation has been applied to the empirical interatomic-distance data for the elementary metals to obtain a nearly complete set of single-bond radii. These radii have been compared with the normal covalent... 

Forty six years ago, on the basis mainly of empirical arguments, I formulated a description of the interatomic forces in metals (2) that had some novel features. I pointed out that according to this view the metallic bond is very closely related to the ordinary covalent (shared-electron-pair) bond some of the electrons in each atom in a metal are involved with those of neighboring atoms in an interaction described as covalent-bond... 

In Fig. 1 there is shown the structure of the orthorhombic crystal iodine as determined with the use of x-rays.2 The atoms are joined in pairs to form molecules by strong shared-electron-pair bonds, and the molecules are... 

Soon after the quantum revolution of the mid 1920s, Linus Pauling and John C. Slater expanded Lewis s localized electronic-structural concepts with the introduction of directed covalency in which bond directionality was achieved by the hybridization of atomic orbitals.1 For normal and hypovalent molecules, Pauling and Slater proposed that sp" hybrid orbitals are involved in forming shared-electron-pair bonds. Time has proven this proposal to be remarkably robust, as has been demonstrated by many examples in Chapter 3. 

The convenient name covalent bond, which we shall often use in this book. in place of the more cumbersome expressions shared-electron-pair bond or electron-pair bond, was introduced by Langmuir (loc. cit. [7], 868). Lewis preferred to include under the name chemical bond a more restricted class of interatomic interactions than that giwri by ov definition f the chemical bond is at all times and in all molecules men ly a pair. i vtror s held jointly by two atoms --Lewis, op. cit. p. 78). 

In a few molecules and crystals it is convenient to describe the interactions between the atoms in terms of the one-electron bond and the three-electron bond. Each of these bbnds is about half as strong as a shared-electron-pair bond each might be described as a half-bond.1 There are also many other molecules and crystals with structures that may be described as involving fractional bonds that result, from the resonance of bonds between two or more positions. Moat of these molecules and crystals have a smaller number of valence electrons than of stable bond orbitals. Substances of this type are called electron-deficient substances. The principal types of electron-deficient substances are discussed in the following sections (and in the next chapter, on metals). 

Lewis s concept of shared electron pair bonds allows for four-electron double bonds and six-electron triple bonds. Carbon dioxide (C02) has two carbon-oxygen double bonds, and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon-nitrogen triple bond. 

The relationship between atomic radii, ionic radii, van der Waals radii, and the electride-ion model of a shared-electron-pair bond is shown in Fig. 29. M represents a relatively large atomic core of an electroposi-... 

Thus the hydrogen atom, with one electron, can achieve the helium structure by taking up another electron, to form the hydride anion, H . But the hydrogen atom can also achieve the helium structure by sharing its electron with the electron of another hydrogen atom, to form a shared-electron-pair bond. Each of the two atoms thus contributes one electron to the shared electron pair. The shared electron... 

In this molecule, the hydrogen fluoride molecule, there is a single covalent bond (shared-electron-pair bond), which holds the hydrogen atom and the fluorine atom firmly together. The distance between the nuclei of these two atoms is 0.92A, according to experimental determination made by the study of the spectrum of the gas. 

Shared-electron pair bond (covalent bond). 

In 1916 the formula with shared-electron-pair bonds was proposed by the American chemist Gilbert Newton Lewis. The large dielectric constant of water was interpreted as showing that the water molecule is not linear, but is bent, with the angle between the two bonds estimated as approximately 110°. 

The boron-boron bonds in this crystal are not ordinary bonds. Boron, with Z = 5 and two electrons in its inner shell, has only three valence electrons, and could form only three shared-electron-pair bonds. But most of the atoms form six bonds. We conclude that these bonds are not ordinary covalent bonds, but are bonds of another kind, involving only one electron per bond. They are called half-bonds a half-bond is only about one-half as strong as an ordinary covalent bond, which involves two electrons (rather than one) shared between two atoms. 

In this addition compound, the N to B bond is referred to as a dative covalent bond, as although the two-electron pair bond is indistinguishable for a normal shared electron pair bond, the source of the electron pair is from one of the atoms of the bond and not equally from both. 

This completes the Lewis structures of CH and CCI4 in which each C has an octet with four bonding pairs, each H has an configuration involving a shared electron pair, and each of the four Cl atoms has an octet, involving one shared electron pair (bonding) and three unshared electron pairs (long pairs). 

This completes the Lewis structures of PCI5 and SF respectively, in which the P has a ten electron core involving five shared electron pair bonds, and S has a twelve electron core involving six shared electron pair bonds. Neither the central P or S atoms has a lone pair of electrons. 

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