Welcome
Content
Introduction
Linear Combination of Atomic Orbitals
H2 and He2 - The Simplest Examples
The Effect of Electronegativity
p-Orbitals and p-Overlap
Combining AOs to Build MOs
The AH2 Walsh Diagram
Larger Molecules
Reactions
Elementary Symmetry
Glossary
Molecules in VRML
General Information - Installation - Use
| 2.2. The Effect of Electronegativity Electronegativity measures the tendency of an atom to accept electrons. Electronegative elements tend to accept electrons from less electronegative ( more electropositive) ones. Thus, electronegative atoms are usually negatively charged (they have more electrons than they need to equalise the core charge) and, oddly enough, electropositive ones are positively charged. In orbital terms, the AO energy levels of electronegative elements lie lower in energy than those of electropositive ones. Figure 2.4 gives the Pauling electronegativities and the electronic configuration for neutral atoms. Figure 2.4 Pauling electronegativities (a) and schematic electronic configuration (b) (images courtesy of Web Elements, Dr Mark Winter, Sheffield). The effect of differing electronegativities on the MOs of the simple two s-orbital system shown above is shown in Figure 2.5 for LiH. Figure 2.5 Orbital interaction diagram for LiH, illustrating the effect of differing electronegativities. The AOs of H and Li have been drawn equally large to emphasise the contributions to the MOs. The s-MO, which is closer to the energy level of the s-AO of the more electronegative hydrogen, is composed of a larger proportion of this AO than of the higher energy lithium s-AO. In contrast, the situation is reversed for the s*-MO, which is closer in energy to the s-AO of the more electropositive lithium. This situation can be expressed in the LCAO expression for the two MOs: Ys = c1 Y s(H) + c 2Y s(Li) Ys* = c1 Y s(H) - c 2Y s(Li) where Ys and Ys* are the two molecular orbitals, Y s(X) the s-AO on element X and the coefficients c1 - c4 determine the relative weights of the AOs in the MOs. These coefficients are defined normalised, so that the total probabilty of finding an electron in the orbital is one. In this case this normalisation is by: c12 + c22 = 1 and c32 + c42 = 1 for the s- and s*-MOs, respectively. The coefficients can be used to divide the electrons between the two atoms involved in the orbital. Because the s-MO is doubly occupied, the number of electrons, NH on the hydrogen atom can be defined as This procedure is known as a population analysis, and is often used to assign net atomic charges, which, however, have no physical meaning, but are simply defined to help interpret the electronic nature of the compound to be studied. It follows that the more electronegative atom, which has the larger coefficient in the s-MO, also has the more negative net atomic charge. Demonstration 2 illustrates the effect of changing the relative electronegativities of the two elements shown in Figure 2.5 on the shapes of the MOs, the bonding energy, the dipole moment and the net atomic charges. H2 and He2 -The Simplest Examples The Effect of Electronegativity Next: p-Orbitals and p-Overlap Combining AOs to Build MOs The AH2 Walsh Diagram