Bonding in diatomic molecules

The basic principles dealing with the molecular orbital description of the bonding in diatomic molecules have been presented in the previous section. However, somewhat different considerations are involved when second-row elements are involved in the bonding because of the differences between s and p orbitals. When the orbitals being combined are p orbitals, the lobes can combine in such a way that the overlap is symmetric around the intemuclear axis. Overlap in this way gives rise to a a bond. This type of overlap involves p orbitals for which the overlap is essentially "end on" as shown in Figure 3.5. For reasons that will become clear later, it will be assumed that the pz orbital is the one used in this type of combination. 

The valence bond method is older and follows quite naturally the notion of two atoms combining to form a molecule by sharing of electrons in atomic orbitals. In this section, we describe bonding in diatomic molecules. 

One of the interesting successes of the molecular orbital approach to bonding in diatomic molecules is the fact that molecules such as 02 and B2 are correctly predicted to be paramagnetic, but the valence bond structures for these molecules are unsatisfactory. Properties for many diatomic species are shown in Table 2.5. 

At this point we switch from consideration of pairs of hybrid orbitals in each bond to bonding and antibonding combinations in each bond. We do this in the way discussed for a simple polar bond in diatomic molecules (Section 1-D) by writing a general linear combination of the two hybrids. 

Table 9 Strengths of bonds in diatomic molecules of the alkaline earth metals with halogens (kcalmoD )...

The VB model constructs wave functions to describe localized electron-pair bonds. The model describes bonding in diatomic molecules, including the... 

In the example of H2, the bond order = (2 - 0) = 1 (a single bond). In diatomic molecules, a bond order of 1 corresponds to a single bond, a bond order of 2 to a double bond, and so forth. 

A relationship between the bonding in diatomic molecules and in solids has been demonstrated. It follows from the so-called Universal Binding Energy Relation (UBER), applicable to both metals and covalent diatomic molecules. The energy and the interatomic separation are scaled in the following way ... 

The nature of chemical bonding in diatomic molecules was extensively studied by Bader and co-workers (Bader, 1963, 1981 Bader et al., 1967a,b Bader and Jones, 1961 Bader and Henneker, 1965 Bader and... 

In Chapter 1, we considered three approaches to the bonding in diatomic molecules ... 

Table 4. Quantum chemical electronegativity and hardness indices for atoms bonded in diatomic molecules. Absolute electronegativity and hardness of molecules, calculated by the same method, is also shown. After Ref. [31]...

Also, in Chapter 3 we introduced the concept of molecular-orbital theory to explain bonding in diatomic molecules. In Chapter 4, we will extend this useful quantum-mechanical concept to polyatomic molecules. In addition, in the final section of Chapter 4, we will examine how molecular shape and bonding affect the interactions of molecules with one another. 

It is relatively simple to assign bond entiialpies to bonds in diatomic molecules. The bond enthalpy is just the energy required to break the diatomic molecule into its component atoms. However, for bonds fliat occur only in polyatomic molecules (such as the C—H bond), we must often utili2e average bond entiialpies. For example, the enthalpy change for tiie following process (called atomization) can be used to define an average bond entiialpy for tiie C—H bond. 

Let us consider the implications of the electrostatic theorem for chemical bonding in diatomic molecules. We take the internuclear axis as the z axis (Fig. 14.4). By symmetry the X and y components of the effective forces on the two nuclei are zero. [Also, one can show that the z force components on nuclei a and b are related by = P z,b (Prob. 14.40). The effective forces on nuclei a and b are equal in magnitude and opposite in direction.]... 

Hybridization of atomic orbitals. To account for the bonding in diatomic molecules like HF or F2, we picture direct overlap of s and/or p orbitals of isolated atoms. But how can we account for the shape of a molecule like methane from the shapes and orientations of C and H atomic orbitals A C atom ([He] 2f2p ) has two valence electrons in the spherical 2s orbital and one each in two of the three mutually perpendicular 2p orbitals. If the half-filled p orbitals overlap the Is orbitals of two H atoms, two C— H bonds would form with a 90° H—C—H bond angle. But methane has the formula CH4, not CH2, and its bond angles are 109.5°. 

Our treatment of molecular orbital theory in this book will be limited to descriptions of bonding in diatomic molecules consisting of elements from the first two periods of the periodic table (H through Ne). 

MOLECULAR BONDING AND STRUCTURE 461 Spinor Bonds in Diatomic Molecules... 

The Heitler-London method is obviously not limited to the treatment of the formation of the bonds in diatomic molecules. In CH4, for example, if we let ci> C2> C3> be the orbitals occupied by four of the electrons of the carbon atom, and hi> h2, be... 

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