Geometry lone pairs effect

This effect of the position of the lone pair is relevant not only in the case of phthalocyaninato ligands, but also can be a clue to the intriguing behaviour of the [Dy(DOTA)] complex, where the rotation of a water molecule changes the magnetic properties [13]. A PCE or an REC model cannot account for the effect of such a rotation, but an LPEC model would predict a dramatic effect, since the change in the position of the lone pair effectively means a completely different geometry. 

In cyclic systems such as 1, the dominant conformation is the one with the maximum anomeric effect. In the case of 1, only conformation lA provides the preferred antiperiplanar geometry for both oxygens. Antiperiplanar relationships are indicated by including lone pairs in the oxygen orbitals. Other effects, such as torsional strain and nonbonded repulsion, contribute to the conformational equilibrium, of course. Normally, a value of about 1.5 kcal/mol is assigned to the stabilization due to an optimum anomeric interaction in an acetal. 

AX5E molecules with one lone pair and AX5=Y molecules have a square pyramidal geometry in which the XAX angles are less than 90°. Some examples are given in Figure 9.21. There is an interesting, but unexplained, difference between the effect of a lone pair... 

Tricoordinate phosphorus is essentially nonplanar in nature and due to the significant s character of the lone pair in a compound such as phosphole, effective overlap with the carbon p orbitals is inhibited and the compound is nonaromatic. Geometry optimizations and aromaticity analyses performed by Glukhovtsev et al. <1996JPC13447> have shown both pentaphosphole 1 and the bicyclic octaphosphane P8 to be both planar and aromatic in character. 

This molecule (type AB3U) has a tetrahedral electronic geometry and a pyramidal molecular geometry. Cl (EN = 3.0) is more electronegative than As (EN = 2.1). The polar As-Cl bond dipoles oppose the effect of the lone pair. The molecule is only slightly polar (Section 8-8). 

This molecule (type AB2U2) has a tetrahedral electronic geometry and an angular molecular geometry. Oxygen (EN = 3.5) is less electronegative than F (EN = 4.0). The O-F bond dipole opposes the effect of the two lone pairs of electrons and so, OF2 is polar (Section 8-9). 

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