Nitrogen molecular orbital theory

Lewis dot diagrams of nitric oxide compared to the nitrosonium ion and molecular nitrogen. Each bond contains one electron from each atom. These simple diagrams fail to properly account for the effective bond order of 2.5 predicted by molecular orbital theory and must be only considered as illustrative. The dimer of two nitric oxide molecules has five bonds, which is the same as two individual molecules. Thus, nitric oxide remains dissociated at room temperatures. [Pg.3]

Most of biological chemistry can be understood in terms of simple ball and stick models. The chemistry of nitric oxide and related oxides is more intimidating because its patterns of bonding depend strongly on quantum mechanics and molecular orbital theory. But the basics can be grasped by comparison to other molecules and a simple consideration of where nitrogen sits in the periodic table. [Pg.18]

From the heat of combustion of benzenamine we know that it has a 3 kcal mole-1 larger stabilization energy than benzene (Table 21-1). This difference in stabilization energies can be ascribed in either valence-bond or molecular-orbital theory to delocalization of the unshared pair of electrons on nitrogen over the benzene ring. The valence-bond structures are... [Pg.1113]

The discussion of molecular structure in terms of Lewis structures invokes the concept of resonance to explain certain physical properties of molecules, such as why two bond lengths are equal. Molecular orbital theory does not use the concept of resonance. Using the azide ion, N3 , as an example, draw the tr-type molecular orbital that leads to bonding between the three nitrogen atoms. [Pg.286]

The triplet state di-TC-methane reactivity of the methanoquinoline systems (61, 62) have been studied. The results of the irradiations are shown in Schemes 3 and 4 where it can be seen that each compound usually affords two products. The results indicate that the pyridine nitrogen does not manifestly alter the photochemical behaviour of the compounds. A detailed analysis of the reasons for the observed regioselectivity is made on the basis of molecular orbital theory. The di-ic-methane reactivity of the benzonorbornadienes (63) still excites considerable interest. Paquette and Burke have studied the triplet state reactivity of the derivatives (63) in an attempt to establish the influence of bridgehead substitution. Thus the irradiation of (63a) results in the formation of both isomers (64a) and <65a) in 42 and 58 It respectively. Similar yields are shown for the irradiation of (63b) when the two products are obtained in similar yields to the above. However, irradiation of the derivative (63c) results in the formation of only the cyclized product (64c) while the derivative with the bridgehead cyano (63d) yields both (64d) and (65d) but in 10 and 90 X yields respectively. The authors present arguments to explain the observed specificities leading to products of either bridgehead control or vinylic control of the biradical... [Pg.256]

Experimental thermodynamic data show that the N2 molecule is stable, is diamagnetic, and has a very high bond energy, 946 kj/mol. This is consistent with molecular orbital theory. Each nitrogen atom has seven electrons, so the diamagnetic N2 molecule has 14 electrons. [Pg.360]

The experimental discovery of Ns" salt stimulated a number of theoretical studies. Bartlett and colleagues [136] studied the stability of salt. It represents a potential solid nitrogen rocket fuel that would be much more efficient than currently rocket propellant. Unfortunately, due to the unavailability of Ns", a direct experimental test of this ion pair is not feasible. The stability of another ion pair, Ns s", has been investigated by Kortus et al. [137]. Unlike Ns s" for which a stability minimum was predicted as two-ion-pair clusters, the Ns" N3" ion pair can spontaneous isomerize to azidopentazole with lower energy and the latter will decay to molecular N2 spontaneously [137]. More recently, using ab initio molecular orbital theory, Dixon et al. [138] predicted that neither the Ns s" nor the N5" 3" ion pair are stable both should decompose spontaneously into N3 radicals and N2. [Pg.212]

These calculations are of special interest because Purcell and coworkers have been interested in the structure-activity relationships of this series of compounds for several years and have compared several approaches to structure-activity relationships with the same compounds. Earlier they (27, 28) reported on detailed studies of the effect of partition coefficients (benzene-water), electric moments, and electronic structures on the relative activity of the congeners. Purcell (28) used Huckel molecular orbital theory to calculate the net charges on all the atoms of each compound. The only apparent success was the apparent correlation of the net charge at the amide nitrogen atom with activity. This was later shown to be caused by the statistically significant correlation between the 7T value and the net charge on the nitrogen atom (29). [Pg.113]

The nitrogen molecule is isoelectronic with carbon monoxide, which we considered earlier. It is best considered on molecular orbital theory. The electronic structure of nitrogen is ls22s22p]c2p[y2pxz. We are dealing with the combination of two such atoms, and we will feed the electrons into the corresponding available molecular orbitals. [Pg.46]

This agrees quite well with simple molecular orbital theory. We assume that the molecule is planar, and that the atoms are joined by a-bonds in sp hybridization. Ten electrons are used up for these bonds, and additionally four are at inner Is shells of the two nitrogen atoms. There are totally seventeen electrons. The remaining three are placed in jr-orbitals. These are constmcted as a linear combination of atomic orbitals according to the LCAO method ... [Pg.16]

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