Lone pair VSEPR theory

The shapes of covalent compounds are determined by the tendency for bonding pairs to be as far apart as possible whilst lone pairs have a greater effect than bonding pairs (VSEPR theory). 

According to VSEPR theory, the most stable arrangement of the three lone pairs of electrons would be in the equatorial position, as shown in (1), where they would be less crowded. Therefore, a linear structure is the correct molecular geometry of the molecule. 

The Cl—F and Cl—Cl bonds in the cation are then formed by the overlap of the half-filled sp3 hybrid orbitals of the central chlorine atom with the half-filled p-orbitals of the terminal Cl and F atoms. Thus, by using sp3 hybridization, we end up with the same bent molecular geometry for the ion as that predicted by VSEPR theory (when the lone pairs on the central atom are ignored)... 

It is noteworthy that the to-bonded structure for ArF6 differs from that predicted by VSEPR theory. ArF6 is predicted to be of octahedral (Oh) symmetry, with three mutually perpendicular F i- Ar -h F triads and an s-type lone pair. In contrast, VSEPR predicts a pentagonal bipyramid (or other seven-vertex polyhedron) with some or all F-Ar-F angles less than 90°. The calculated equilibrium structure is in agreement with the co-bonding model. 

Incidentally, it may be emphasized that (i) the MYKO-XYL structure constitutes a challenge to group theory in that it has no ternary symmetry and that (ii) the structure of MYKO-CgHg exhibits a tricky counter-example to Gillespie s VSEPR model, the lone pairs of the two exocyclic N atoms on each P being strictly parallel.. . ... 

The basic principle underlying VSEPR theory is that valence electron pairs, whether they re lone pairs or they occur within bonds, prefer to be as far from one another as possible. There s no sense in crowding negative charges any more than necessary. 

In general, it is assumed in VSEPR theory that the repulsive interactions between bonding (b) and non-bonding or lone ( ) electron pairs are ordered ... 

Figure 6.19 Examples of VSEPR theory predictions of molecular shapes in a series of hypothetical dehydration reactions in which element E has one lone pair of electrons...

These lowest oxidation states can be associated with a coordination number of two and a linear arrangement of the ligands about the central atom. On VSEPR theory these are 10-electron species, with a trigonal bipyramidal arrangement of three lone pairs of electrons in equatorial positions and two halogen ligands in axial positions. 

The three hexafluorohalogen(VII) cations have been shown to have octahedral symmetry by spectroscopic methods. This undistorted octahedral geometry is predicted by VSEPR theory, since no lone pairs of electrons are involved. 

The octalluoroxenaies are the most stable xenon compounds known they can be heated to 400 °C without decomposition. The anions have square antiprismatic geometry. They, too, present a problem to VSEPR theory analogous to that of XeF6 since they should also have a stereochemically active lone pair of electrons that should lower the symmetry of the anion. If the steric crowding theory is correct, however, the presence of eight ligand atom/, could force the lone pair into a stereochemically inert s Orhital. 

Atoms are bound into molecules by shared pairs of electrons. Electrons dislike each other because like charges repel each other. Therefore, whether they are lone pairs of electrons or bonding pairs of electrons, electron pairs try to get as far apart in space as is geometrically possible. There is a fancy name that summarizes these simple ideas the VSEPR theory, which stands for Valence Shell Electron Pair Repulsion Theory. Even though the VSEPR theory is founded on fundamentally simple ideas, it is a tremendously powerful tool for predicting the shapes of molecules. 

The angle between lone pairs is larger than that between two bonded pairs, in agreement with the VSEPR (valence-shell electron-pair repulsion) theory. 

The structures of these ions normally conform to those predicted by the VSEPR theory, as shown in Fig. 17.2.2. Since the anion XY has two more electrons than the cation XY+, they have very different shapes. The anion IFj- is planar with lone pairs occupying the axial positions of a pentagonal bipyramid. In [Me4N](IF6), IFg is a distorted octahedron (C3V symmetry) with a sterically active lone pair, whereas both BrFg and ClFg are octahedral. The anion IF " has the expected square antiprismatic structure. 

In the vast majority of cases in which six coordination is observed, the bonding can be viewed as arising from the interaction of all three cr -orbitals with a halide anion, i.e., all three in S. Because the three orbitals are all trans to the primary E-X bonds, such a situation leads naturally to octahedral coordination. Moreover, in cases in which the primary and secondary bonds are the same length, i.e., where A = 0 and a three-center, four-electron bonding model is appropriate, a regular octahedron is the result. Such a structure is clearly at odds with simple VSEPR theory, which is predicated on the lone pair(s) occupying specific stereochemical sites, but stereochemical inactivity of the lone pair tends to be the rule rather than the exception in six-coordinate, seven-electron pair systems Ng and Zuckerman (102) have reviewed this topic for p-block compounds in general. 

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