Bonding in methane

What is the average enthalpy of atomization of the four C—H bonds in methane Compare this value with the accepted value of the C—H bond enthalpy. 

Caibon has eight electrons in its valence shell in both methane and carbon tetrafluoride. By forming covalent bonds to four other atoms, carbon achieves a stable electron configuration analogous to neon. Each covalent bond in methane and carbon tetrafluoride is quite strong—comparable to the bond between hydrogens in Fl2 in bond dissociation energy. 

Before we describe the bonding in methane, it is worth pointing out that bonding theories attempt to describe a molecule on the basis of its component atoms bonding theories do not attempt to explain how bonds for-rn. Thus, the world s methane does not come... 

Section 2.6 Bonding in methane is most often described by an orbital hybridization model, which is a modified for m of valence bond theory. Four equivalent 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 I5 orbital gives a a bond. 

As the table indicates, C—H bond dissociation energies in alkanes are approximately 375 to 435 kJ/mol (90-105 kcal/mol). Homolysis of the H—CH3 bond in methane gives methyl radical and requires 435 kJ/mol (104 kcal/mol). The dissociation energy of the H—CH2CH3 bond in ethane, which gives a primary radical, is somewhat less (410 kJ/mol, or 98 kcal/mol) and is consistent with the notion that ethyl radical (primary) is more stable than methyl. 

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. 

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 D values may be easy or difficult to measure, and they can be estimated by various techniques, but there is no question as to what they mean. With E values the matter is not so simple. For methane, the total energy of conversion from CH4 to C + 4 H (at 0 K) is 393 kcal mol (1644 kJ mol ). " Consequently, E for the C—H bond in methane is 98 kcal mol (411 kJ mol ) at OK. The more usual practice, though, is not to measure the heat of atomization (i.e., the energy necessary to convert a compound to its atoms) directly but to calculate it from the heat of combustion. Such a calculation is shown in Figure 1.11. 

There are many other molecules in which some of the electrons are less localized than is implied by a single Lewis structure and can therefore be represented by two or more resonance structures. For example, the three bonds in the carbonate ion all have the same length of 131 pm, which is intermediate between that of the C—O single bond in methanol (143 pm) and that of the C=0 double bond in methanal (acetaldehyde) (121 pm). So the carbonate ion can be conveniently represented by the following three resonance structures ... 

The dissociation energy for C-H bond in methane (E = 436 kj/mol) is one of the highest among all organic compounds. Its electronic structure (i.e., the lack of n- and n-electrons), lack of polarity, and any functional group makes it extremely difficult to thermally decompose the methane molecule into its constituent elements. 

To calculate the average bond enthalpy of the (C-H) bond in methane, E (C-H), the average of the four bond dissociation enthalpies is calculated ... 

The reaction enthalpy and thus the RSE will be negative for all radicals, which are more stable than the methyl radical. Equation 1 describes nothing else but the difference in the bond dissociation energies (BDE) of CH3 - H and R - H, but avoids most of the technical complications involved in the determination of absolute BDEs. It can thus be expected that even moderately accurate theoretical methods give reasonable RSE values, while this is not so for the prediction of absolute BDEs. In principle, the isodesmic reaction described in Eq. 1 lends itself to all types of carbon-centered radicals. However, the error compensation responsible for the success of isodesmic equations becomes less effective with increasingly different electronic characteristics of the C - H bond in methane and the R - H bond. As a consequence the stability of a-radicals located at sp2 hybridized carbon atoms may best be described relative to the vinyl radical 3 and ethylene 4 ... 

This value of the C-H bond enthalpy does not correspond to the enthalpy of dissociation of the carbon-hydrogen bond in methane, as represented in Equation (4.35). 

The idea here is just the same, except for inevitable refinements and details, for the formation of aU covalent bonds. So the basic ideas for chemical bonding in methane, ammonia, water, and so on, are the same. 

Now we can consider the bonding in methane. Using orbital overlap as in the hydrogen molecule as a model, each sp orbital of carbon can now overlap with a 1 orbital of a hydrogen atom, generating a bonding molecular orbital, i.e. a ct bond. Four such... 

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