Molecular orbitals Resonance structures

We see, then, that resonance theory gives us exactly the same picture of the allyl radical that we obtained from molecular orbital theory. Structure C describes the carbon-carbon bonds of the allyl radical as partial double bonds. The resonance structures A and B also tell us that the unpaired electron is associated only with the Cl and C3 atoms. We indicate this in structure C by placing a 8 beside Cl and C3. Because resonance structures A and B are equivalent, the electron density from the unpaired electron is shared equally by Cl and C3. 

Molecular orbitals are useful tools for identifying reactive sites in a molecule. For example, the positive charge in allyl cation is delocalized over the two terminal carbon atoms, and both atoms can act as electron acceptors. This is normally shown using two resonance structures, but a more compact way to see this is to look at the shape of the ion s LUMO (the LUMO is a molecule s electron-acceptor orbital). Allyl cation s LUMO appear s as four surfaces. Two surfaces are positioned near- each of the terminal car bon atoms, and they identify allyl cation s electron-acceptor sites. 

Unlike the stable molecule N2O, the sulfur analogue N2S decomposes above 160 K. In the vapour phase N2S has been detected by high-resolution mass spectrometry. The IR spectrum is dominated by a very strong band at 2040 cm [v(NN)]. The first ionization potential has been determined by photoelectron spectroscopy to be 10.6 eV. " These data indicate that N2S resembles diazomethane, CH2N2, rather than N2O. It decomposes to give N2 and diatomic sulfur, S2, and, hence, elemental sulfur, rather than monoatomic sulfur. Ab initio molecular orbital calculations of bond lengths and bond energies for linear N2S indicate that the resonance structure N =N -S is dominant. 

Computational approaches remain almost untouched, and the main reason probably is the inevitable structural complication which is required of molecules and ions to gain sufficient resonance stabilization. Recently, a molecular orbital calculation has been applied to a simplified hydrocarbon salt model by Fujimoto et al. (1991). 

Our treatment of O2 shows that the extra complexity of the molecular orbital approach explains features that a simpler description of bonding cannot explain. The Lewis structure of O2 does not reveal its two unpaired electrons, but an MO approach does. The simple (t-tt description of the double bond in O2 does not predict that the bond in 2 is stronger than that in O2, but an MO approach does. As we show in the following sections, the molecular orbital model has even greater advantages in explaining bonding when Lewis structures show the presence of resonance. 

If the molecule contains multiple bonds, construct the n bonding system using molecular orbital theory, as described in this section and in the remaining pages of Chapter 10. Watch for resonance structures, which signal the presence of delocalized electrons. 

Triatomic species can be linear, like CO2, or bent, like O3. The principles of orbital overlap do not depend on the identity of the atoms involved, so all second-row triatomic species with 16 valence electrons have the same bonding scheme as CO2 and are linear. For example, dinitrogen oxide (N2 O) has 16 valence electrons, so it has an orbital configuration identical to that of CO2. Each molecule is linear with an inner atom whose steric number is 2. As in CO2, the bonding framework of N2 O can be represented with sp hybrid orbitals. Both molecules have two perpendicular sets of three tt molecular orbitals. The resonance structures of N2 O, described... 

Resonant forms of molecules are more stable than the structures from which they form. The new orbitals extend over the entire molecule. This allows the electrons to have longer wavelengths and correspondingly lower energy. Delocalization also plays a role in the last two topics covered in this chapter molecular orbitals and metallic bonding. 

In a series of studies of the spectroscopy and photochemistry of nickel(O) -a-diimine complexes, the structural differences among the complexes NiL2 and Ni(CO)2L (L Q-diimine) have been examined by means of molecular orbital calculations and electronic absorption Raman resonance studies.2471, 472 Summing up earlier work, the noninnocence of a-diimine ligands with a flat — N=C—C=N— skeleton in low-valent Ni chemistry and the course of substitution reactions of Ni° complexes with 1,4-diaza-1,3-dienes or a,a -bipyridine have been reviewed.2473... 

It is readily apparent that the three a bonds are capable of holding the six bonding electrons in the a t and e molecular orbitals. The possibility of some 7r-bonding is seen in the molecular orbital diagram as a result of the availability of the a2" orbital, and in fact there is some experimental evidence for this type of interaction. The sum of the covalent radii of boron and fluorine atoms is about 152 pm (1.52 A), but the experimental B-F bond distance in BF3 is about 129.5 pm (1.295 A). Part of this "bond shortening" may be due to partial double bonds resulting from the 7r-bonding. A way to show this is by means of the three resonance structures of the valence bond type that can be shown as... 

Various theoretical methods (self-consistent field molecular orbital (SCF-MO) modified neglect of diatomic overlap (MNDO), complete neglect of differential overlap (CNDO/2), intermediate neglect of differential overlap/screened approximation (INDO/S), and STO-3G ab initio) have been used to calculate the electron distribution, structural parameters, dipole moments, ionization potentials, and data relating to ultraviolet (UV), nuclear magnetic resonance (NMR), nuclear quadrupole resonance (NQR), photoelectron (PE), and microwave spectra of 1,3,4-oxadiazole and its derivatives <1984CHEC(6)427, 1996CHEC-II(4)268>. 

A variety of acyclic and cyclic S-N compounds decompose at moderate temperatures (100-150 °C) with the formal loss of a symmetrical NSN fragment, but this molecule has never been detected. The lowest energy isomer, linear NNS, is generated by flash vacuum pyrolysis of 5-phenyl-l,2,3,4-thia-triazole.40 Ab initio molecular orbital calculations indicate that the resonance structure N = N+-S is dominant.41... 

The electronic structure of dichorodiphenylplumbane was calculated by the SCF-MS (self-consistent field multiple scattering) molecular orbital method and compared to that of dichlorodiphenylstannane. The results suggest that one has to look for 35C1 NQR (nuclear quadrupole resonance) frequencies of dichorodiphenylplumbane in the 5-6 MHz region1613. 

A In this exercise we will combine valence-bond and molecular orbital methods to describe the bonding in the SO3 molecule. By invoking a 7t-bonding scheme, we can replace the three resonance structures for SO3 (shown below) with just one structure that exhibits both cr-bonding and delocalized n--bonding. 

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