Bonding description

Although AILs (bond paths) in AIM theory often coincide with our basic understanding of what a chemical bond is, this is not always the case. There has been a lot of debate on this topic in the literature, with the main source of disagreement arising from insufficient distinction between the AIM concept and our 

Examples of molecular graphs for some representative molecules. The nuclear positions are drawn as open circles, the bond critical points (BCP) as filled circles. 

Contour map of p at the BCP perpendicular to the N—N AIL. The positions of the hydrogen atoms in front of and behind the plane indicate the molecular orientation. 

---

FIGURE 3 10 Bent bonds in cyclopropane (a) The orbitals involved in carbon-carbon bond formation overlap in a region that is displaced from the internuclear axis (b) The three areas of greatest negative electrostatic potential (red) correspond to those predicted by the bent bond description... 

In keeping with the "bent-bond" description of Figure 3.10, the carbon-carbon bond distance in cyclopropane (151 pm) is slightly shorter than that of ethane (153 pm) and cyclohexane (154 pm). The calculated values from molecular models (see Learning By Modeling) reproduce these experimental values. 

Figure 5-20. Reaction coordinate diagram generated from a valence bond description of initial and final state configurations.

This valence bond description leads to an interesting conclusion. Because the transition state occurs at the point where the initial and final state VB configurations cross, the transition state receives equal contributions from each. This is so whether the transition state is early or late. Thus, the nucleophile Y and the leaving group X possess about equal charge densities in the transition state. This conclusion means that an early transition state is not (in this sense) reactantlike , for a reactantlike transition state should have most of the charge on Y. Similarly, a late transition state is not necessarily productlike. This view is at variance with other interpretations. 

It follows from the preceding discussion that the unbranched H bond can be regarded as a 3-centre 4-electron bond A-H B in which the 2 pairs of electrons involved are the bond pair in A-H and the lone pair on B. The degree of charge separation on bond formation will depend on the nature of the proton-donor group AH and the Lewis base B. The relation between this 3-centre bond formalism and the 3-centre bond descriptions frequently used for boranes, polyhalides and compounds of xenon is particularly instructive and is elaborated in... 

The extent to which each hybrid is incorporated into the full bonding description of the molecule will depend on the extent to which 3d orbitals... 

Cooper, D. L., Gerratt, J., and Raimondi, M. The Spin-Coupled Valence Bond Description of Benzenoid Aromatic Molecules. 153, 41-56 (1990). 

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. 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

The Mulliken charges of the iron and sulfur atoms, and especially the bond index between the bridging sulfurs (see Tables 1 and 2), favor a bonding description of a single S - S bond that stems from the [Fe - - Fe] core... 

The Zintl-Klemm concept evolved from the seminal ideas of E. ZintI that explained the structural behavior of main-group (s-p) binary intermetaUics in terms of the presence of both ionic and covalent parts in their bonding description [31, 37]. Instead of using Hume-Rother/s idea of a valence electron concentration, ZintI proposed an electron transfer from the electropositive to the electronegative partner (ionic part) and related the anionic substructure to known isoelectronic elemental structures (covalent part), e.g., TK in NaTl is isoelectro-nic with C, Si and Ge, and consequenUy a diamond substructure is formed. ZintI hypothesized that the structures of this class of intermetallics would be salt-like [16b, 31 f, 37e]. 

Our work described in this section clearly illustrates the importance of the nature of the cations (size, charges, electronegativities), electronegativity differences, electronic factors, and matrix effects in the structural preferences of polar intermetallics. Interplay of these crucial factors lead to important structural adaptations and deformations. We anticipate exploratory synthesis studies along the ZintI border will further result in the discovery of novel crystal structures and unique chemical bonding descriptions. 

We begin with the Lewis structure of PH3, and apply the principles of orbital overlap. The bonding description of the molecule must be consistent with the bond angles, so we must pay particular attention to the orientation of the valence orbitals. 

As with the aiuminum atom in triethyiaiuminum, the bonding description of the zinc atom in dimethylzinc inciudes vacant p orbitais as weii as hybrid orbitais for bonding. As Figure 10-16 shows, a set of sp hybrid orbitals makes use of only one valence p orbital. The remaining two p orbitals are perpendicular to each other and perpendicular to the pair of hybrids. 

The carbon and nitrogen atoms have a p orbital left over after construction of the. s p hybrids. These two p orbitals are perpendicular to the plane that contains the five nuclei. Side-by-side overlap of the p orbitals gives a jrbond that completes the bonding description of this molecule ... 

Sum the electrons allocated in Steps 1-3. The result must match the total number of valence electrons used in the Lewis structure, hi addition, a complete bonding description must account for all the valence orbitals. 

All these properties of metals are consistent with a bonding description that places the valence electrons in delocalized orbitals. This section describes the band theory of solids, an extension of the delocalized orbital ideas... 

---