Carbon valence orbitals hybridization

The carbon atom with its 2s2, 2p2 electron configuration can form a maximum of four bonds which can be either single, double or triple. As a consequence as three main connectivities result depicted in Fig. 1. The valence orbitals hybridize forming, in the elemental state, sp, sp2 or sp3 hybrid orbitals. The data in Fig. 1 illustrate that the local bonding geometry and the carbon-carbon bond distances vary significantly with the connectivity as the bond order (and bond strength) increases from the sp3 to the sp hybridization state. 

Heimann, R.B., Evsyukov, S.E., and Koga, Y. (1997). Carbon aUotropes a suggested classification scheme based on valence orbital hybridization. Carbon, 35, 1654-8. 

The overlap integrals between the carbon valence orbitals and the hydrogen Is orbital decreases in the order of HCsCH, CH2=CH2, and CH4 (with the hybridization index for the carbon atom equal to 1, 2, and 3, respectively). Assuming that the C-H bond strength increases with the overlap integral, we can conclude that the overlap integrals determined in (i) correlate well with the given spectral... 

Figure 1.3 Pictorial representation of taking a linear combination of the four carbon valence orbitals to form sp -hybrid orbitals.

Figure 1.4 Result of hybridization of all four of the carbon valence orbitals.

Section 2 6 Bonding m methane is most often described by an orbital hybridization model which is a modified form of valence bond theory Four equiva lent sp hybrid orbitals of carbon are generated by mixing the 2s 2p 2py and 2p orbitals Overlap of each half filled sp hybrid orbital with a half filled hydrogen Is orbital gives a ct bond... 

The properties of tert butyl cation can be understood by focusing on its structure which IS shown m Figure 4 9 With only six valence electrons which are distributed among three coplanar ct bonds the positively charged carbon is sp hybridized The unhybridized 2p orbital that remains on the positively charged carbon contains no elec Irons Its axis is perpendicular to the plane of the bonds connecting that carbon to the three methyl groups... 

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

The valence-bond concept of orbital hybridization described in the previous four sections is not limited to carbon compounds. Covalent bonds formed by-other elements can also be described using hybrid orbitals. Look, for instance, at the nitrogen atom in methylamine, CH3NH2, an organic derivative of ammonia (NH3) and the substance responsible for the odor of rotting fish. 

Divalent carbon species called carbenes are capable of fleeting existence. For example, methylene, CH2, is the simplest carbene. The two unshared electrons in methylene can be either spin-paired in a single orbital or unpaired in different orbitals. Predict the type of hybridization you expect carbon to adopt in singlet (spin-paired) methylene and triplet (spin-unpaired) methylene. Draw a picture of each, and identify the valence orbitals on carbon. 

We saw in Chapter 1 that the carbon-carbon double bond can be described in two ways. In valence bond language (Section 1.8), the carbons are sp2-hybridized and have three equivalent hybrid orbitals that lie in a plane at angles of 120° to one another. The carbons form a cr bond by head-on overlap of sp2 orbitals and a tt bond by sideways overlap of unhybridized p orbitals oriented... 

A great deal of evidence has shown that carbocations are planar. The divalent carbon is 5p2-hybridized, and the three substituents are oriented to the corners of an equilateral triangle, as indicated in Figure 6.9. Because there are only six valence electrons on carbon and all six are used in the three a bonds, the p orbital extending above and below the plane is unoccupied. 

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. 

The carbon-oxygen double bond of a carbonyl group is similar in many respects to the carbon-carbon double bond of an alkene. The carbonyl carbon atom is s/ 2-hybridized and forms three valence electron remains in a carbon p orbital and forms a tt bond to oxygen by overlap with an oxygen p orbital. The oxygen atom also has two nonbonding pairs of electrons, w hich occupy its remaining two orbitals. 

It may be proper at this stage to lead the reader back to the stage where we constructed the localized orbitals of a CH2 group. At that time two valence orbitals were set aside—the 2pv orbital, and the outer (2s, 2pr) hybrid. Both of these orbitals lie in the. r, y plane. Now in our description of cyclopropane, we used bond orbitals to describe the CC bonding these bond orbitals are derived from in-plane (xy y) hybrids on each carbon. The two hybrids which are required on each carbon atom—in ordet to participate in two bond orbitals—are built precisely from the 2py orbital and the (2s, 2pj.) out combination on each CH2 group. 

The hybridized orbital approach is a simplified way of predicting the geometry of a molecule by mixing the valence orbitals of its atoms. For example, methane (CH ) is composed of a carbon atom with an electron configuration of Is 2s 2p. The hydrogen atom has an electron configuration of Is. The geometry of the methane... 

Cyclic organic compounds as a basic variant of carbon nanostructures. Apparently, not only inner-atom hybridization of valence orbitals of carbon atom takes place in cyclic structures, but also total hybridization of all cycle atoms. 

Turning now to direct theoretical evaluations, we consider Ar as the displacement (on the C=C axis) of the centroid of the tt orbital with respect to the center 1. Of course, such a displacement can differ from zero only if some hybridization is allowed, which, in the case of a tt orbital, must consist in admixture of the suitable dTT orbital. The hybrid in question was determined [222] from 4-3IG calculations with d polarization functions for carbon and optimization of all scale factors, followed by a calculation of in situ valence orbitals, and of their characteristics, according to Del Re and Barbier [143]. The inward shifts (on the C=C axis) of the tt orbital centroids are close to 0.03 A (Table 11.1). 

The VSEPR theory assumes that the four electrons from the valence shell of the carbon atom plus the valency electrons from the four hydrogen atoms form four identical electron pairs which, at minimum repulsion, give the observed tetrahedral shape. To rationalize the tetrahedral disposition of four bond-pair orbitals with those of the 2s and three 2p atomic orbitals of the carbon atom, sp3 hybridization is invoked. 

Answer to 4(d). We can consider the orbitals of cycloheptatrienylidene to arise from the interaction of the orbitals of hexatriene and the valence orbitals of a di-coordinated carbon atom (a 2p orbital and an sp" hybrid orbital). The orbitals of hexatriene may be obtained from an SHMO calculation. The interaction diagram is shown in Figure B7.2. The p orbital of the carbene site is raised as a result of the dominant interaction with 713 of hexatriene. The orbital 714, which is closest in energy to the carbene s p orbital, does not interact because of symmetry, and 5 interacts less strongly because the coefficients at the terminal positions of the hexatriene are... 

We are now ready to account for the bonding in a methane molecule. An unpaired electron occupies each of carbon s sp3 hybrid orbitals. Each of these four electrons can pair with an electron in a hydrogen ls-orbital. Their overlapping orbitals form four cr-bonds (Fig. 3.17). Because the hybrid orbitals point toward the corners of a tetrahedron, so do the a-bonds they form. All four bonds are identical because they are formed from the same blend of atomic orbitals. The valence-bond description is now consistent with experiment. 

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