Carbon atom hybridized orbitals

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...

Hirshfeld (1964) pointed out that bond bending not only occurs in ring systems, but also results from steric repulsions between two atoms two bonds apart, referred to as 1-3 interactions. The effect is illustrated in Fig. 12.3. The atoms labeled A and A are displaced from the orbital axes, indicated by the broken lines, because of 1-3 repulsion. As a result, the bonds defined by the orbital axes are bent inwards relative to the internuclear vectors. When one of the substituents is a methyl group, as in methanol [Fig. 12.3(b)], the methyl-carbon-atom hybrid reorients such as to maximize overlap in the X—C bond. This results in noncolinearity of the X—C internuclear vector and the three-fold symmetry axis of the methyl group. Structural evidence for such bond bending in acyclic molecules is abundant. Similarly, in phenols such as p-nitrophenol (Hirshfeld... 

The bond angles of the carbon atoms in benzene are 120°. All carbon atoms are s p -hybridized, and each carbon atom has a single unhybridized p orbital perpendicular to the plane of the ring. The carbon p -hybridized orbitals overlap to form the ring of the benzene molecule. Because the C—C bond lengths are 1.39 A, the p orbitals are close enough to overlap efficiently and equally all round the ring. 

The relative importance of a and r contributions to the overall bonding is unclear, but several different combinations of relative strengths lead to limiting case models. When there are 2 electrons in the forward (T-bond and 2 electrons in the ir-backbond, there are 2 bonding electrons for each metal-carbon bond. This is mathematically equivalent to 2tr-bonds and a metallocyclopropane structure (72). This model does not necessitate strict sp3 hybridization at the carbon atoms. Molecular orbital calculations for cyclopropane (15) indicate that the C—C bonds have higher carbon atom p character than do the C—H bonds. Thus, the metallocyclopropane model allows it interactions with substituent groups on the olefin (68). 

Ketene looks pretty unlikely It is CH2=C=0 with two tt bonds (C=C and C=0) to the same carbon atom. The orbitals for these jt bonds must be orthogonal because the central carbon atom is sp hybridized with two linear o bonds and two p orbitals at right angles both to the o bonds and to each other. Can such a molecule exist When acetone vapour is heated to very high temperatures (700-750 °C) methane is given off and ketene is supposed to be the other product. What is isolated is a ketene dimer (C4H4O2) and even the structure of this is in doubt as two reasonable structures can be written. 

This information may be interpreted in terms of the unpaired electron being confined to a carbon sp hybrid orbital (4ai) built up by carbon 2s and 2p and oxygen 2p atomic orbitals. The 2pj character of the 4ai molecular orbital can be estimated by comparison with the integral ... 

There are now eight different spatial orbitals, hybrid orbitals, the other four being close to atomic hydrogen s-orbitals. The expansion of each of the VB orbitals in terras of M the basis functions located on the nuclei allows the orbitals to distort from the pure atomic shape. The SCVB wave... 

One of the simplest functional groups consists of a single halogen atom, which we take to be chlorine for illustrative purposes. The chlorine atom forms a a bond to a carbon atom by overlap of its 3p orbital with a hybridized orbital on the carbon. The hybridized orbital may be sp, sp, or sp depending on the bonding in the hydrocarbon frame. Alkyl halides form when mixtures of alkanes and halogens (except iodine) are heated or exposed to light. 

More recently Newton, Schulman and Manus (41) considered the relationship between experimental /(C-C) values in 12 different hydrocarbons and the product of the percent s character in the two bonding carbon atomic hybrids obtained from localization of INDO molecular orbitals (see Fig. 1). They found that... 

Similar arguments apply to the acetylene molecule. Here we must assume that in each carbon atom hybridization takes place between the zs orbital and only one of the 2p orbitals (say px), giving rise to two collinear sp or cr bonds. The unaltered py and pz orbitals stand perpendicular to these bonds and to each other. In the molecule the cr bonds will be responsible for the bonding H-C-C-H, but again there... 

If we consider the carbon atom to be sp hybridised from 2s+2px orbitals, then we could have a linear O—C—O framework containing a bonds, by overlap of the two carbon sp hybrid orbitals with the two oxygen 2px orbitals. This is illustrated in Figure 20a. 

Fig. 8.23. The ethane molecule in the antiperiplanar configuration (a). The localized oibilal of the CH bond (b) and the localized orbital of the CC bond (c). The carbon atom hybrid forming the CH bond is quite similar to the hybrid forming the CC bond.

The indicated bond in 1,3-butadiene is formed by the overlap of two sp -hybridized carbons, while the indicated bond in 1-butene is formed by the overlap of one sp -hybridized and one sp -hybridized carbon. sp -Hybridized orbitals are smaller because they have greater s character (an s orbital holds its electrons closer to the nucleus of an atom), so the bond made from two sp hybrids is shorter. 

In order to form a monocarbide, the valence electrons of the carbon atom hybridize with the spd band of the metal atom. It is likely that the metal orbitals are the dP-sj hybridization since the typical octahedral grouping of the metal atoms centered on the carbon atom has six bonds to the six comers of the octahedron, thus favoring the M-C bond. Indeed, the cP-sjP hybridization is common in the Group IV metals (Ti, Zr,... 

As you see from the orbital diagram for the hybridized C atom, a single 2p orbital still remains on each carbon atom. These orbitals are perpendicular to the plane of the hybrid orbitals that is, they are perpendicular to the —CH2 plane. Note that the two —CH2 planes can rotate about the carbon-carbon axis without affecting the overlap of the hybrid orbitals. As these planes rotate, the 2p orbitals also rotate. When the —CH2 planes rotate so that the 2p orbitals become parallel, the orbitals overlap to give a IT bond (Figure 10.265). 

The sigma bond framework (in green) forms by the overlap of sp hybrid orbitals on each carbon atom H atom Is orbitals overlap these hybrid orbitals to form C—H bonds. Delocalized pi orbitals form by the overlap of carbon atom 2p orbitals that are perpendicular to the plane of the molecule they give orbitals with lobes above and below this plane. Only the lowest-energy pi orbital is shown (yellow). 

FIGURE 6.12 The formation of the triple bond in the organic compound acetylene (C2H2). (a) The orbitals on each carbon atom hybridize as in Figure 6.11, and (b) the other open slots are filled as shown. 

Fig. 10.7 The 2s and three 2p orbitals of a carbon atom hybridize, and the resulting hybrid orbitals point toward the corners of a regular tetrahedron.

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