Ammonia, bonding molecular orbitals

The tetramethylammonium salt [Me4N][NSO] is obtained by cation exchange between M[NSO] (M = Rb, Cs) and tetramethylammonium chloride in liquid ammonia. An X-ray structural determination reveals approximately equal bond lengths of 1.43 and 1.44 A for the S-N and S-O bonds, respectively, and a bond angle characteristic bands in the IR spectrum at ca. 1270-1280, 985-1000 and 505-530 cm , corresponding to o(S-N), o(S-O) and (5(NSO), respectively. Ab initio molecular orbital calculations, including a correlation energy correction, indicate that the [NSO] anion is more stable than the isomer [SNO] by at least 9.1 kcal mol . ... 

The final molecule of this series is methane, the tetrahedral structure of which follows if a fourth unit positive charge is removed from the nucleus in the ammonia lone-pair direction. There are now four equivalent bonding orbitals, which may be represented approximately as linear combinations of carbon s-p hybrid and hydrogen Is functions. The transformation from molecular orbitals into equivalent orbitals or vice versa is exactly the same as for the neon atom. 

But how do we account for the bond angles in water (104°) and ammonia (107°) when the only atomic orbitals are at 90° to each other All the covalent compounds of elements in the row Li to Ne raise this difficulty. Water (H2O) and ammonia (NH3) have angles between their bonds that are roughly tetrahedral and methane (CH4) is exactly tetrahedral but how can the atomic orbitals combine to rationalize this shape The carbon atom has electrons only in the first and second shells, and the Is orbital is too low in energy to contribute to any molecular orbitals, which leaves only the 2s and 2p orbitals. The problem is that the 2p orbitals are at right angles to each other and methane does not have any 90° bonds. (So don t draw any either Remember Chapter 2.). Let us consider exactly where the atoms are in methane and see if we can combine the AOs in such a way as to make satisfactory molecular orbitals. 

This order is well illustrated by the various angles in sulfur diflu-oride in Figure 3-39 as determined by ab initio molecular orbital calculations [92], This is also why, for example, the bond angles H-N-H of ammonia, 106.7°, are smaller than the ideal tetrahedral value, 109.5°. Unless stated otherwise, the parameters in the present discussion are taken from the Landolt Bornstein Tables [93],... 

Yeo, G. A. and Ford, T. A.,Ab initio molecular orbital calculations of the infrared spectra of hydrogen bonded complexes of water, ammonia, and hydroxylamine. Part 6. The infrared spectrum of the water-ammonia complex. Can. J. Chem. 69,632-637 (1991). 

Turi, L., and Dannenberg, J.J., Molecular orbital studies of the nitromethane-ammonia complex. An unusually strong C—H "N hydrogen bond, J. Phys. Chem. 99, 639-641 (1995). 

The directionality in the bonding between a d-block metal ion and attached groups such as ammonia or chloride can now be understood in terms of the directional quality of the d orbitals. In 1929, Bethe described the crystal field theory (CFT) model to account for the spectroscopic properties of transition metal ions in crystals. Later, in the 1950s, this theory formed the basis of a widely used bonding model for molecular transition metal compounds. The CFT ionic bonding model has since been superseded by ligand field theory (LFT) and the molecular orbital (MO) theory, which make allowance for covalency in the bonding to the metal ion. However, CFT is still widely used as it provides a simple conceptual model which explains many of the properties of transition metal ions. 

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