Nitrogen valence, orbital

The C2v character table shows that none of the nitrogen valence orbitals transforms as a2. However, px transforms as bt and therefore can participate in tt bonding. 

Ammonia fits the same pattern. Bonding in NH3 uses all the nitrogen valence orbitals, so the hybrids are sp, including one 5 orbital and all three p orbitals, with overall tetrahedral symmetry. The predicted HNH angle is 109.5°, narrowed to the actual 106.6° by repulsion from the lone pair, which also occupies an sp orbital. 

The nitrogen 2 s and 2p valence AO s conveniently oriented for interaction, are shown in Fig. 5. This orientation naturally divides the valence AO s into three orbital sets, namely, the 2s set, the 2 px set and the 2 py, 2 pz set. The latter two sets we refer to as radial and tangential p orbitals, respectively. The radial orbitals point toward the molecular center and the tangential orbitals are perpendicular to the radial orbitals. Let us consider these sets separately and imagine various GO s placed one by one between the nitrogen atoms. For each GO we try to find that combination of nitrogen orbitals which has the same symmetry, until we have as many symmetry orbitals (SO s) as there are nitrogen valence orbitals in the set. 

Both UPS and XPS of solids are useful techniques. So far as studies of adsorption by surfaces are concerned we would expect UPS, involving only valence orbitals, to be more sensitive. For example, if we wish to determine whether nitrogen molecules are adsorbed onto an iron surface with the axis of the molecule perpendicular or parallel to the surface it would seem that the valence orbitals would be most affected. This is generally the case but, because ultraviolet photoelectron spectra of solids are considerably broadened, it is the X-ray photoelectron spectra that are usually the most informative. 

For the same reason we discussed for oxygen atoms, the nitrogen atom is most stable when it has the maximum number of partially filled valence orbitals. This keeps the electrons as far apart as possible. The most stable state of the nitrogen atom is as follows ... 

The nonbonding pair on the nitrogen atom occupies the 2 s orbital, which Is the only valence orbital of this atom not used for bonding. 

A ligand must have a lone pair of electrons that it donates to form a bond to the metal. For example, the Ni— bonds in [Ni (NH3)g form by overlap of an empty valence orbital on the metal with the lone pair. s orbital on the nitrogen atom. The water ligands in [Ni (H2 0) ] " coordinate to the metal in a similar manner, with the oxygen atoms donating lone pairs to form Ni—O bonds. 

Ammonia is a prime example of a Lewis base. In addition to its three N—H bonds, this molecule has a lone pair of electrons on its nitrogen atom, as Figure 21-1 shows. Although all of the valence orbitals of the nitrogen atom in NH3 are occupied, the nonbonding pair can form a fourth covalent bond with a bonding partner that has a vacant valence orbital available. 

The simplest type of Lewis acid-base reaction is the combination of a Lewis acid and a Lewis base to form a compound called an adduct. The reaction of ammonia and trimethyl boron is an example. A new bond forms between boron and nitrogen, with both electrons supplied by the lone pair of ammonia (see Figure 21-21. Forming an adduct with ammonia allows boron to use all of its valence orbitals to form covalent bonds. As this occurs, the geometry about the boron atom changes from trigonal planar to tetrahedral, and the hybrid description of the boron valence orbitals changes from s p lo s p ... 

Nitrogen has one filled and three half filled valence orbitals. Two nitrogen atoms form three bonds with their three half-filled orbitals. The remaining free pairs of electrons (one on each N atom) are placed around the nitrogen atoms. 

It is well known that a change in the H—N—H valence angle of NH3 is important for the energy of the mn orbital. In the transition from the pyramidal to planar conformation, this orbital destabilizes appreciably with decreasing contribution of the nitrogen 2s orbital. This is also reflected in the very low ionization potentials of planar amines (see below). 

One final point about covalent bonds involves the origin of the bonding electron pair. Although most covalent bonds form when two atoms each contribute one electron, bonds can also form when one atom donates both electrons (a lone pair) to another atom that has a vacant valence orbital. The ammonium ion (NH4+), for example, forms when the two lone-pair electrons from the nitrogen atom of ammonia, NH3, bond to H +. Such bonds are called coordinate covalent bonds. 

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