Bonding, normal covalent

Both these molecules exist in the gaseous state and both are trigonal planar as indicated by reference to Table 2.8. However, in each, a further covalent bond can be formed, in which both electrons of the shared pair are provided by one atom, not one from each as in normal covalent bonding. For example, monomeric aluminium chloride and ammonia form a stable compound ... 

The dotted line shows the second bond formed by hydrogen, the bond called the hydrogen bond. It is usually dotted to indicate that it is much weaker than a normal covalent bond. Consideration of the boiling points in Figure 17-14, on the other hand, shows that the interaction must be much stronger than van der Waals forces. Experiments show that most hydrogen bonds release between 3 kcal/mole and 10 kcal/mole upon formation ... 

C O , that is, the normal covalent carbon-oxygen double bond,- the estimated bond energy 6.60 v.e., then the ketones would show a resonance energy of 1.11 v.e. arising from the + —... 

It is shown that the energy of a normal covalent bond A-B between unlike atoms is probably represented more closely by the geometric mean of the bond energies for A-A and B-B than by their arithmetic mean. 

Extreme Ionic Bonds and Normal Covalent Bonds.—Before discussing the nature of actual bonds it is desirable to specify the sense in which the terms ionic and covalent will be used. 

The Additivity of the Energies of Normal Covalent Bonds. The Hydrogen Halides and the Halogen Halides.—It is found that there exists a convincing body of empirical evidence in support of the postulate6 that the energies of normal covalent bonds are additive that is... 

It is perhaps desirable to point out that the bond type has no direct connection with ease of electrolytic dissociation in aqueous solution. Thus the nearly normal covalent molecule HI ionizes completely in water, whereas the largely ionic HF is only partially ionized. 

Bond Energies and the Relative Electronegativity of Atoms.—In Table II there are collected the energies of single bonds obtained in the preceding sections. One additional value, obtained by a method to be described later, is also included 1.44 v. e. for N N. Under each bond energy is given the value for a normal covalent bond, calculated from additivity, and below that the difference A. It is seen that A is positive in twenty of the twenty-one cases. The exception, C I, may be due to experimental error, and be not real. 

The quantum mechanical argument used in deriving the original electronegativity scale involved the amount of ionic character of a normal covalent bond A—B, and it was evident that the amount of ionic character and accordingly the value of the electric dipole moment of the bond would be closely correlated with the difference Ax = xA — xB of the two atoms A and B. In the first edition of The Nature of the Chemical Bond (1939) the following equation was advanced ... 

It is customary to use the line between two atomic symbols, A—B, to represent a normal covalent bond, with the usual amount of ionic character. In the discussion that follows, A B is used to represent a pure covalent single bond, and A+B to represent the ionic structure that is hybridized with A B to give A—B. A pure covalent double bond is represented by A=B. Thus in the molecule NF3 we might describe each NF bond as involving N+ F , N F , and N F +, that is, as having some covalent double-bond character. 

Bismuth is intermediate in the transition from a metallic to a normal covalent structure each atom shows the effect of its normal tricovalence by having three nearest neighbors, at 3.10 A. and it has also three near neighbors at the larger distance 3.47 A. The respective bond... 

The values of f (l) given in the table for electronegative atoms are their normal covalent single-bond radii28 (except for boron, discussed below). The possibility that the radius 0.74 A. of Schomaker and Stevenson29 should be used for nitrogen in the metallic nitrides should be borne in mind. 

After rising at copper and zinc, the curve of metallic radii approaches those of the normal covalent radii and tetrahedral covalent radii (which themselves differ for arsenic, selenium, and bromine because of the difference in character of the bond orbitals, which approximate p orbitals for normal covalent bonds and sp3 orbitals for tetrahedral bonds). The bond orbitals for gallium are expected to be composed of 0.22 d orbital, one s orbital, and 2.22 p orbitals, and hence to be only slightly stronger than tetrahedral bonds, as is indicated by the fact that R(l) is smaller than the tetrahedral radius. 

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... 

The normal valence (normal covalence) of sulfur, corresponding to its position in the periodic table, is 2. The electronic structure of the normal bicovalent sulfur atom is the argononic structure with bond orbitals about... 

The selective reaction of K with the a-rh boron is explained by the metastable character of this variety at high T and by the existence in the structure of weak, three-center B—B bonds that correspond to particularly long B — B distances (2.03 X 10 pm a normal covalent B—B bond is 1.76 X 10 pm). 

Immobilization by chemical bonding gives strong, irreversible attachments to a solid support. The bonds are normally covalent but they can be electrostatic. Typical supports are functionalized glass and ceramic beads and fibers. Enzymes are sometimes cross-linked to form a gel. Occasionally, enz5anes can be flocculated while retaining catalytic activity. 

Coupling between fluorine and a hydrogen, a carbon or another fluorine that may be separated by many bonds (four, five, six or more) can result from overlap of electronic orbitals occupied by lone pair electrons which are unshared and therefore not involved in normal covalent bonding. The term applied to this effect, through space is somewhat misleading, since all isotropic coupling must be transmitted in some way by electrons, either in bonds or in unshared pairs. 

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