Methane orbital bonds

At this stage, it looks as though electron promotion should result in two different types of bonds in methane, one bond from the overlap of a hydrogen ls-orbital and a carbon 2s-orbital, and three more bonds from the overlap of hydrogen Is-orbitals with each of the three carbon 2/ -orbitals. The overlap with the 2p-orbitals should result in three cr-bonds at 90° to one another. However, this arrangement is inconsistent with the known tetrahedral structure of methane with four equivalent bonds. 

Hybridization of the carbon to which a proton is attached also influences electron density. As the proportion of s character increases from sp to sp to sp orbitals, bonding electrons move closer to carbon and away from the protons, which then become deshielded. For this reason, methane and ethane resonate at 8 0.23 and 0.86, respectively, but ethene resonates at 8 5.28. Ethyne (acetylene) is an exception in this regard, as we shall see. Hybridization contributes to shifts in strained molecules, such as cyclobutane (8 1.98) and cubane (8 4.00), for which hybridization is intermediate between sp and sp. ... 

The functional group of a thiol is an —SH (sulfhydryl) group bonded to an sp hybridized carbon. Figure 10.4 shows a Lewis structure and a ball-and-stick model of methanethiol, CHjSH, the simplest thiol. The C—S—H bond angle in methane-thiol is 100.3°. By way of comparison, the H—S—H bond angle in H2S is 93.3°. If a sulfur atom were bonded to two other atoms by fully hybridized sp hybrid orbitals, bond angles about sulfur would be approximately 109.5°. If, instead, a sulfur atom were bonded to two other atoms by unhybridized 3p orbitals, bond angles would be... 

It was stated earlier that CHa prefers bond angles of 120°, and methane prefers bond angles of 109.5°. How do we achieve such bond angles when the s and p atomic orbitals are not oriented at these angles The s orbitals are spherical and so have no directionality in space, and the p orbitals are oriented at 90° angles with respect to each other. We need a conceptual approach to understand how s and p atomic orbitals can accommodate these experimentally determined molecular bond angles. The most common approach is the idea of hybridization, first introduced by Pauling. 

FIGURE 2.9 A structure for methane, CH4. Bonds are formed by the overlap of four hydrogen Ir orbitals wth the 2p and 2s atomic orbitals of one carbon atom. This model requires some 90° bond angles and requires the bonds to be of different lengths. Carbon s li electrons are shown in parentheses because they are not used in bonding. 

Table 10 Fluorinated methanes mean bond energies MBE) of GEL, and CF4 and C-F thermo-chemical bond energies (TBE) of CEL,. F , = 1 to 3 and net atomic charges obtained by natural atomic orbital analysis (NAO) of calculated electron densities [65]...

Hence we have two molecular orbitals, one along the line of centres, the other as two sausage-like clouds, called the n orbital or n bond (and the two electrons in it, the n electrons). The double bond is shorter than a single C—C bond because of the double overlap but the n electron cloud is easily attacked by other atoms, hence the reactivity of ethene compared with methane or ethane. 

A vexing puzzle m the early days of valence bond theory concerned the fact that methane is CH4 and that the four bonds to carbon are directed toward the corners of a tetrahedron Valence bond theory is based on the overlap of half filled orbitals of the connected atoms but with an electron configuration of s 2s 2p 2py carbon has only two half filled orbitals (Figure 2 8a) How can it have bonds to four hydrogens ... 

The axes of the sp orbitals point toward the corners of a tetrahedron Therefore sp hybridization of carbon is consistent with the tetrahedral structure of methane Each C—H bond is a ct bond m which a half filled Is orbital of hydrogen over laps with a half filled sp orbital of carbon along a line drawn between them... 

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... 

Fig. 1.3. Valenee bond stmetural representation of methane resulting from overlap of H li orbitals with four equivalent sp orbitals of earbon. ...

Fig. 1.20. Atomic orbital combinations giving rise to bonding molecular orbitals for methane.

For a molecule as simple as Fl2, it is hard to see much difference between the valence bond and molecular orbital methods. The most important differences appear- in molecules with more than two atoms. In those cases, the valence bond method continues to view a molecule as a collection of bonds between connected atoms. The molecular- orbital method, however, leads to a picture in which the sane electron can be associated with many, or even all, of the atoms in a molecule. We ll have more to say about the similarities and differences in valence bond and molecular- orbital theory as we continue to develop their principles, beginning with the simplest alkanes methane, ethane, and propane. 

The same kind of orbital hybridization that accounts for the methane structure also accounts for the bonding together of carbon atoms into chains and rings to make possible many millions of organic compounds. Ethane, C2H6, is the simplest molecule containing a carbon-carbon bond. 

Pentadienyl radical, 240 Perturbation theory, 11, 46 Propane, 16, 165 n-Propyi anion conformation, 34 n-Propyl cation, 48, 163 rotational barrier, 34 Propylene, 16, 139 Protonated methane, 72 Pyrazine, 266 orbital ordering, 30 through-bond interactions, 27 Pyridine, 263 Pyrrole, 231... 

When we try to apply VB theory to methane we run into difficulties. A carbon atom has the configuration [HeJ2s22pvl2p l,1 with four valence electrons (34). However, two valence electrons are already paired and only the two half-filled 2/ -orbitals appear to be available for bonding. It looks as though a carbon atom should have a valence of 2 and form two perpendicular bonds, but in fact it almost always has a valence of 4 (it is commonly tetravalent ) and in CH4 has a tetrahedral arrangement of bonds. 

We are now ready to account for the bonding in methane. In the promoted, hybridized atom each of the electrons in the four sp3 hybrid orbitals can pair with an electron in a hydrogen ls-orbital. Their overlapping orbitals form four o-bonds that point toward the corners of a tetrahedron (Fig. 3.14). The valence-bond description is now consistent with experimental data on molecular geometry. 

FIGURE 3.14 Each C H bond in methane is formed by the pairing of an electron in a hydrogen U-orbital and an electron in one of the four sp hybrid orbitals of carbon. Therefore, valence-bond theory predicts four equivalent cr-bonds in a tetrahedral arrangement, which is consistent with experimental results. 

The development of molecular orbital theory (MO theory) in the late 1920s overcame these difficulties. It explains why the electron pair is so important for bond formation and predicts that oxygen is paramagnetic. It accommodates electron-deficient compounds such as the boranes just as naturally as it deals with methane and water. Furthermore, molecular orbital theory can be extended to account for the structures and properties of metals and semiconductors. It can also be used to account for the electronic spectra of molecules, which arise when an electron makes a transition from an occupied molecular orbital to a vacant molecular orbital. 

There are four equivalent orbitals, each called sp, which point to the comers of a regular tetrahedron (Fig. 1.4). The bond angles of methane (CH4) would thus be expected to be 109° 28, which is the angle for a regular tetrahedron. 

Although the hybrid orbitals discussed in this section satisfactorily account for most of the physical and chemical properties of the molecules involved, it is necessary to point out that the sp orbitals, for example, stem from only one possible approximate solution of the Schrddinger equation. The i and the three p atomic orbitals can also be combined in many other equally valid ways. As we shall see on page 12, the four C—H bonds of methane do not always behave as if they are equivalent. 

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