Triple bonds molecular orbital theory

Molecular orbital theory may provide an explanation for stereochemical differences between carboxylate-metal ion and phosphate-metal ion interactions. Detailed ab initio calculations demonstrate that the semipo-lar 1 0 double bond of RsP=0 is electronically different from the C=0 double bond, for example, as found in H2C=0 (Kutzelnigg, 1977 Wallmeier and Kutzelnigg, 1979). The P=0 double bond is best described as a partial triple bond, that is, as one full a bond and two mutually perpendicular half-7r bonds (formed by backbonding between the electrons of oxygen and the empty d orbitals of phosphorus). Given this situation, a lone electron pair should be oriented on oxygen nearly opposite the P=0 bond, and these molecular orbital considerations for P=0 may extend to the phosphinyl monoanion 0-P=0. If this extension is valid, then the electronic structure of 0-P=0 should not favor bidentate metal complexation by phosphate this is in accord with the results by Alexander et al. (1990). 

This case study shows that CO molecules do not have significant vibrational energy unless the temperature is quite high. This happens because CO has a triple bond and, therefore, a large force constant k = 1902 N m ). The correlation between force constant and bond order in diatomic molecules is explained by molecular orbital theory, and is summarized in Figure 6.20. Other diatomic molecules will behave differently, as determined by their structure and the Boltzmann distribution. 

When carbon vaporizes at extremely high temperatures, among the species present in the vapor is the diatomic molecule C2. Write a Lewis formula for C2. Does your Lewis formula of C2 obey the octet rule (C2 does not contain a quadruple bond.) Does C2 contain a single, a double, or a triple bond Is it paramagnetic or diamagnetic Show how molecular orbital theory can be used to predict the answers to questions left unanswered by valence bond theory. 

Historically, the application of molecular orbital theory to the electronic structures of isoelec-tronic 14-electron molecules such as acetylene, HCN, N2, and O2 was an excellent pioneering demonstration of the value of quantum chemistry. Within the framework of molecular orbital theory, the C - C bond in acetylene is a triple bond involving one a-bond, and two orthogonal 7t-bonds. The a-bond is formed by two sp-hybrid orbitals from each carbon, and the two 71-bonds are formed from the perpendicular p-orbitals. Alternatively, the so-called bent or banana bonds have been invoked to describe the multiple C-C bonds in acetylene (Fig. 1-1) [3-5]. This creates a conceptual dilemma, though one bonding model can be transformed to the other by appropriate linear combinations. It is now realized that both... 

The theories of HLPS might be called electron-pairing theories if Lewis is called an electron-pair theory. It should also be pointed out that the HLPS electron pair differs considerably from Lewis conception of the electron-pair bond, in that the electrons are much less closely associated in this respect it approaches the truth much more closely than does Lewis conception - Pauling and Slater consider a double bond to be merely two ordinary single bonds sticking out from each atom in different directions, and treat the triple bond in a similar way. In this way they do not agree very well with Lewis, nor do they agree with the results obtained from molecular orbital theory. 

Hybridization in Molecules Containing Double and Triple Bonds 337 Molecular Orbital Theory 340... 

How does each of the three major bonding theories (the Lewis model, valence bond theory, and molecular orbital theory) define a single chemical bond A double bond A triple bond How are these definitions similar How are they... 

Valence bond theory does agree fairly well with molecular orbital (MO) theory for homonuclear diatomic molecules that can obey the octet rule H2 (single bond, bond order = 1), Li2 (single bond, bond order = 1), N2 (triple bond, bond order = 3), 02 (double bond, bond order = 2), F2 (single bond, bond order = 1). However, for those molecules that don t, it is more difficult to know if they exist or not and what bond orders they have. MO theory allows us to predict that He2, Be2 and Ne2 do not exist since they have bond orders = 0, and that B2 has bond order = 1 and C2 has bond order = 2. 

The main result that emerges from the discussions of particular eases is that it has proved possible to give a description of a molecule in terms of equivalent orbitals which are approximately localised, but which can be-transformed into delocalised molecular orbitals without any change in the value of the total wave function. The equivalent orbitals are closely associated with the interpretation of a chemical bond in the theory, for, in a saturated molecule, the equivalent orbitals are mainly localised about two atoms, or correspond to lone-pair electrons. Double and triple bonds in molecules such as ethylene and acetylene are represented as bent single bonds, although the rather less localised o-n description is equally valid. 

A two-dimensional Httckel molecular orbital (HMO) theory approach to acetylenic systems yielded n-bond orders of P = 0.894 for the central C—C bond and P = 1.788 for the C C triple bonds in 1,3-butadiyne (209, 210). For comparison, P = 1 for ethylene and P = 2 for acetylene. A different criterion for determining the relative strengths of chemical bonds was used by Politzer and Ranganathan (17). Starting from STO-3G geometries and force constants, they calculated a bond order of 1.34 for the central C-C bond in diacetylene. This corresponds to a bond dissociation energy of 150 kcal/mol [211], which compares with bond orders and bond dissociation energies of 1.14 and 88 kcal/mol for ethane and 1.85 and 163 kcal/mol for ethylene. 

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