Electron orbitals hybrid

The water molecule is composed of two hydrogen atoms covalently bonded to an oxygen atom with tetrahedral (sp3) electron orbital hybridization. As a result, two lobes of the oxygen sp3 orbital contain pairs of unshared electrons, giving rise to a dipole in the molecule as a whole. The presence of an electric dipole in the water molecule allows it to solvate charged ions because the water dipoles can orient to form energetically favorable electrostatic interactions with charged ions. 

We 11 expand our picture of bonding by introducing two approaches that grew out of the idea that electrons can be described as waves—the valence bond and molecular orbital models In particular one aspect of the valence bond model called orbital hybridization, will be emphasized... 

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

The structure of ethylene and the orbital hybridization model for its double bond were presented m Section 2 20 and are briefly reviewed m Figure 5 1 Ethylene is planar each carbon is sp hybridized and the double bond is considered to have a a component and a TT component The ct component arises from overlap of sp hybrid orbitals along a line connecting the two carbons the tt component via a side by side overlap of two p orbitals Regions of high electron density attributed to the tt electrons appear above and below the plane of the molecule and are clearly evident m the electrostatic potential map Most of the reactions of ethylene and other alkenes involve these electrons... 

According to the orbital hybridization model benzene has six tt elec Irons which are shared by all six sp hybridized carbons Regions of high TT electron density are located above and below the plane of the ring... 

An orbital hybridization description of bonding m methylamme is shown m Figure 22 2 Nitrogen and carbon are both sp hybridized and are joined by a ct bond The unshared electron pair on nitrogen occupies an sp hybridized orbital This lone parr IS involved m reactions m which amines act as bases or nucleophiles The graphic that opened this chapter is an electrostatic potential map that clearly shows the concentration of electron density at nitrogen m methylamme... 

What accounts for the stability of conjugated dienes According to valence bond theory (Sections 1.5 and 1.8), the stability is due to orbital hybridization. Typical C—C bonds like those in alkanes result from a overlap of 5p3 orbitals on both carbons. In a conjugated diene, however, the central C—C bond results from conjugated diene results in part from the greater amount of s character in the orbitals forming the C-C bond. 

Fukui applied the orbital mixing rule [1,2, 59] to the orbital hybridization or the deformation of the LUMO of cyclohexanone to explain the origin of the Jt-facial selectivity in the reduction of cyclohexanone. Cieplak [60] proposed that electron delocalization occurs from the bonds into the o orbital of the incipient bonds at the transition state. 

To truly understand the geometry of bonds, we need to understand the geometry of these three different hybridization states. The hybridization state of an atom describes the type of hybridized atomic orbitals (ip, sp, or sp) that contain the valence electrons. Each hybridized orbital can be used either to form a bond with another atom or to hold a lone pair. 

Hybrids of the type sp3 are unjustified for disilane. An important conclusion from the above hybridization statement No. 4 is concerned with the contrasting structures of the radicals SiH3 and CH3. The planar geometry of the methyl radical can readily be explained by the (bond-strengthening) sp2-hy-bridization, while the pyramidal silyl radical is thought to be stabilized (with respect to the planar arrangement) through the s-admixture to the lone electron orbital. 

The orbitals containing the bonding electrons are hybrids formed by the addition of the wave functions of the s-, p-, d-, and f- types (the additions are subject to the normalization and orthogonalization conditions). Formation of the hybrid orbitals occurs in selected symmetric directions and causes the hybrids to extend like arms on the otherwise spherical atoms. These arms overlap with similar arms on other atoms. The greater the overlap, the stronger the bonds (Pauling, 1963). 

The prototype FeCr sigma phase is of particular interest because the free atoms have very nearly the same size (ratio = 1.01), but they condense into a rather intricate structure. In the pure metals, the diameter of Cr is 2.50 A, while that of Fe is 2.48 A. (a difference of less than one percent), and both are bcc. Therefore, the existence of the sigma phase is determined by spd-hybridization of the electron orbitals. It is sometimes called a size-effect phase, but this is not really descriptive. 

The Ni octahedra derive their stability from the interactions of s, p, and d electron orbitals to form octahedral sp3d2 hybrids. When these are sheared by dislocation motion this strong bonding is destroyed, and the octahedral symmetry is lost. Therefore, the overall (0°K) energy barrier to dislocation motion is about COCi/47r where = octahedral shear stiffness = [3C44 (Cu - Ci2)]/ [4C44 + (Cu - C12)] = 50.8 GPa (Prikhodko et al., 1998), and the barrier = 4.04 GPa. The octahedral shear stiffness is small compared with the primary stiffnesses C44 = 118 GPa, and (Cn - C12)/2 = 79 GPa. Thus elastic as well as plastic shear is easier on this plane than on either the (100), or the (110) planes. 

Figure 1.12 The hypothetical formation of methane from an sp -hybridized carbon atom. In orbital hybridization we combine orbitals, not electrons. The electrons can then be placed in the hybrid orbitals as necessary for bond formation, but always in accordance with the Pauli principle of no more than two electrons (with opposite spin) in each orbital. In this illustration we have placed one electron...

It seems to be realistic to relate catalytic activity to the most stable [111] plane of fee metals. Bond (135) describes the electron structure of the this plane. So-called 2g electron orbitals point toward those interstices where metal atoms in the subsequent overlayer would be accommodated. These orbitals have metallic character. So-called orbitals point toward the next nearest neighbor. These are localized and able to form real covalent bonds. The degree of hybridization of these orbitals is imknown. Knor (136) assumes that only orbitals would stick out of the plane, but they are almost completely hybridized. He assumes that the /2g electrons are parts of the electron gas of the metal. The and sites are by no means equivalent. 

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