Molecular structure Trigonal planar molecules

The Lewis formula for the molecule (type AB3) predicts 3 electron groups around the central N atom. Only 1 of the two resonance structures is shown. The electronic and molecular geometries are the same, trigonal planar, because there are no lone pairs of electrons on the N atom (Section 8-6). 

When a molecule is relatively small and/or belongs to a point group of relatively high symmetry, it is possible to elucidate the molecular structure by using the symmetry selection rules discussed in Section 1.14. Molecules of XY2 (linear or bent C2v), XY3 (planar D3h or pyramidal C3v), XY4 (square-planar D4h or tetrahedral Td) and XY5 (trigonal-bipyramidal D3h or tetragonal-pyramidal C4v) types may take one of the structures indicated in parentheses. Since the number of IR/Raman-active vibrations is different for each structure, the most probable structure can be chosen by comparing the number of observed IR/Raman bands with that predicted for each structure by symmetry selection rules. 

To determine the molecular structure, we must count the electron pairs around the sulfur atom. In each resonance structure the sulfur has one lone pair, one pair in a single bond, and one double bond. Counting the double bond as one pair yields three effective pairs around the sulfur. According to Table 13.8, a trigonal planar arrangement is required, yielding a V-shaped molecule ... 

The crystal structures of glycylaminomethylphosphonic acid, and of methane-, ethane-, and propane-diphosphonic acids, have been determined. The unit cell of the last acid contains two molecules, with different conformations. The molecular structures of the constrained phosphite (161), the phosphate (162), and the thio-phosphate (163) have been compared. The nitrogen in the last compound is very nearly trigonal planar, and the large P—N distance (313 pm) shows that there is little P - - N interaction. The phosphazene (164) adopts a novel conformation, ... 

T2.6 The planar from of NHj molecule would have a trigonal planar molecular geometiy with the lone pair residing in nitrogen s p orbital that is perpendicular to the plane of the molecule. This structure belongs to a Djh point group (for discussion of molecular symmetry and point groups refer to Chapter 6). Now we can consult Resource Section... 

For molecular species with other than linear, trigonal planar or tetrahedral-based structures, it is usual to involve d orbitals within valence bond theory. We shall see later that this is not necessarily the case within molecular orbital theory. We shall also see in Chapters 14 and 15 that the bonding in so-called hypervalent compounds such as PF5 and SFg, can be described without invoking the use of J-orbitals. One should therefore be cautious about using sp"d hybridization schemes in compounds of />-block elements with apparently expanded octets around the central atom. Real molecules do not have to conform to simple theories of valence, nor must they conform to the sp"d" schemes that we consider in this book. Nevertheless, it is convenient to visualize the bonding in molecules in terms of a range of simple hybridization schemes. 

Give one example of a compound having a linear molecular structure that has an overall dipole moment (is polar) and one example that does not have an overall dipole moment (is nonpolar). Do the same for molecules that have trigonal planar and tetrahedral molecular structures. 

This method is widely used for elucidation of the molecular structure of inorganic, organic, and coordination compounds. In Part 11 and Appendix HI, the number of infrared- and Raman-active fundamentals is compared for XY3 (planar, Da, and pyramidal, Cj,), XY4 (square-jJanar, D4/ and tetrahedral, Td), XYj (trigonal-bipyramidal, Dj, and tetragonal-pyramidal, 4 ) and other molecules. Recently, the structures of various metal carbonyl compounds (Sec. III-16) have been determined by this simple technique. 

The molecular structure of sulfur dioxide is angular, while the structure of the trioxide is trigonal planar. See Fig. 20.2. It is noteworthy that the bond distances in SO2 and SO3 are 5 or 6 pm shorter than in the monoxide, while the mean bond energies in SO2 and SO3 are respectively 3% greater and 9% smaller than the dissociation energy of the monoxide. These observations indicate that the SO bonds in these molecules are best described as double. 

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