Atomic orbital energies and electronegativity

In spite of the widespread use of electronegativity as a unifying concept in organic chemistry, the electronegativity of an element is almost never included in the periodic table. Redressing this deficiency, Allen strikingly showed his electronegativity scale as the third dimension of the periodic table, and his vivid picture is adapted here as Fig. 1.55. 

As usual, we can tackle the problem with or without using the concept of hybridisation. The C—X bond in a molecule such as methyl chloride, like the C C bond in ethane (Fig. 1.22), has several orbitals contributing 

The important thing for the moment is the comparison between the C C orbitals and the corresponding C Cl orbitals. What we learn about the properties of C—Cl bonds by looking at this one orbital will be the same as we would have learned, at much greater length, from the set as a whole. Alternatively, we can use an interaction diagram for an sp3 hybrid on carbon and an sp3 hybrid on chlorine, and compare the result with 

On account of the loss of symmetry, the chlorine atom has a larger share of the total electron population. In other words, the coefficient on chlorine for the bonding orbital, crCci is larger than that on carbon. It follows from the requirement that the sum of the squares of all the c-values on any one atom in all the molecular orbitals must equal one, that the coefficients in the corresponding antibonding orbital, r ccl must reverse this situation the one on carbon will have to be larger than the one on chlorine. 

C—Cl bond strength represented by these numbers comes from the purely covalent bonding given by 2EC in Fig. 1.58. The other part of the strength of the C Cl bond comes from the electrostatic attraction between the high electron population on the chlorine atom and the relatively exposed carbon nucleus. 

As usual, we can tackle the problem with or without using the concept of hybridisation. The C—X bond in a molecule such as methyl chloride, like the C C bond in ethane, has several orbitals contributing to the force which keeps the two atoms bonded to each other but, just as we could abstract one pair of atomic orbitals of ethane and make a typical interaction diagram for it, so can we now take the corresponding pair of orbitals from the set making up a C—Cl a bond. 

In making a covalent bond between carbon and chlorine from the 2px orbital on carbon and the 3px orbital on chlorine, we have an interaction (Fig. 1.45) between orbitals of unequal energy (-10.7 eV for C and -13.7 eV for Cl, from Table 1.1). The interaction diagram could equally have been drawn using sp3 hybrids on 

We might be tempted at this stage to say that we have a weaker bond than we had for a CC bond, but we must be careful in defining what we mean by a weaker bond. Tables of bond strengths give the CCl bond a strength of 

352 kJ mol-1 (84 kcal moP1), whereas a typical C—C bond strength is a little lower at 347 kJ moP1 (83 kcal mol-1). The C—Cl bond is strong, if we try to break it homolytically to get a pair of radicals, and the C—C bond is easier to break this way. This is what the numbers 352 and 347 kJ mol-1 refer to. 

A picture of the electron distribution in the a orbitals between carbon and chlorine is revealed in the wire-mesh diagrams for methyl chloride in Fig. 1.47, which shows one contour of the major orbital crccl contributing to C—Cl bonding together with the LUMO, r ccl, which incidentally provides the cover illustration for this book. Comparing these with the schematic version in Fig. 1.46, we can see how the back lobe on carbon in rccl includes overlap with the s orbitals on the hydrogen atoms, and that the front lobe in cr ccl wraps back a little behind the carbon atom to include some overlap to the s orbitals of the hydrogen atoms. 

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