Nitrogen atom, orbital overlap

But how can we have generated three basis functions for a doubly degenerate representation The answer is that eqs. (6), (7), and (8) are not LI. So we look for two linear combinations that are LI and will overlap with the nitrogen atom orbitals px and... 

These interactions are most commonly observed for divalent chalcogen atoms and the nitrogen atom (the electron donor D) lies within the X-E-Y (E = S, Se, Te) plane, preferably along the extension of one of the covalent bonds as in 15.3. This anisotropy is a clear indication that these short E N contacts have some bonding character, i.e., they are subject to the geometric restrictions of orbital overlap. Eor example, in the diselenide 15.4 the nitrogen lone pairs are clearly oriented towards the Se-Se linkage. ... 

In contrast with amines, amides (RCONH ) are nonbasic. Amides don t undergo substantial protonation by aqueous acids, and they are poor nucleophiles. The main reason for this difference in basicity between amines and amides is that an amide is stabilized by delocalization of the nitrogen lone-pair electrons through orbital overlap with the carbonyl group. In resonance terms, amides are more stable and less reactive than amines because they are hybrids of two resonance forms. This amide resonance stabilization is lost when the nitrogen atom is protonated, so protonation is disfavored. Electrostatic potential maps show clearly the decreased electron density on the amide nitrogen. 

In the literature discussing these results, the coincidence of the NN bond lengths in diazonium ions with that in dinitrogen seems always to be regarded with complete satisfaction. In the opinion of the present author this close coincidence is somewhat surprising, firstly because of the fact that in diazonium ions one of the nitrogen atoms is bonded to another atom in addition to the N(2) atom, and secondly because work on dual substituent parameter evaluations of dediazoniation rates of substituted benzenediazonium ions clearly demonstrates that the nx orbitals of the N(l) nitrogen atom overlap with the aromatic 7t-electron system (see Sec. 8.4). 

We encounter a different type of bond in a nitrogen molecule, N2. There is a single electron in each of the three 2p-orbitals on each atom (33). However, when we try to pair them and form three bonds, only one of the three orbitals on each atom can overlap end to end to form a (T-bond (Fig. 3.10). Two of the 2/7-orbitals on each atom (2px and 2py) are perpendicular to the internuclear axis, and each one contains an unpaired electron (Fig. 3.11, top). When the electrons in one of these p-orbitals on each N atom pair, the orbitals can overlap only in a side-by-side arrangement. This overlap results in a TT-bond, a bond in which the two electrons lie in two lobes, one on each side of the internuclear axis (Fig. 3.11, bottom). More formally, a 7T-bond has a single nodal plane containing the internuclear axis. Although a TT-bond has electron density on each side of the internuclear axis, it is only one bond, with the electron cloud in the form of two lobes, just as a p-orbital is one orbital with two lobes. In a molecule with two Tr-bonds, such as N2, the... 

The carbon and nitrogen atoms have a p orbital left over after construction of the. s p hybrids. These two p orbitals are perpendicular to the plane that contains the five nuclei. Side-by-side overlap of the p orbitals gives a jrbond that completes the bonding description of this molecule ... 

The triple bond includes two itt bonds formed by the pair-wise overlap of the p orbitals that remain on the carbon and nitrogen atoms. Two figures from different perspectives help illustrate the triple bond ... 

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. 

When we consider double bonds to oxygen, as in carbonyl groups (C=0) or to nitrogen, as in imine functions (C=N), we find that experimental data are best accommodated by the premise that these atoms are sp hybridized (Figure 2.22). This effectively follows the pattern for carbon-carbon double bonds (see Section 2.6.2). The double bond is again a combination of a ct bond plus a jt bond resulting from overlap of p atomic orbitals. The carbonyl... 

Heteroatoms a to a C —II insertion site can activate that site for insertion. It was shown that an acyclic amide such as 2-diazo-lV,jV-diisopropyl-3-oxobutanamide can cyclize smoothly to the /Mactam56. It may be that in this case overlap between the filled orbital on the nitrogen atom and the C-H orbital can increase the electron density in the latter, thus making it more reactive. It may be pertinent that as the starting acyclic amide becomes more product-like, amide resonance decreases, and the electron-donating ability of the nitrogen atom could increase. 

If the nitrogen atom s orbitals in a pyridine-like compound are assumed to be sp2 hybridized, like carbon orbitals, the live 2s and 2p electrons obviously must be allotted to the individual orbitals as shown in Fig. 1. Overlap of the 2sp2 hybrid orbitals of the heteroatom with similar orbitals of the neighboring atoms results in the formation of (T-bonds overlap of its 2pz orbital with the 2pz orbitals of other... 

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