Bonding models, comparison

Reactions of UCI4 with [Li RC(NCy)2 (THF)]2 (R = Me, Bu ) in THF gave the tris(amidinate) compounds [RC(NCy)2]3UCl that could be reduced with lithium powder in THF to the dark-green homoleptic uranium(lll) complexes [RC(NCy)2]3U. Comparison of the crystal structure of [MeC(NCy)2]3U with those of the lanthanide analog showed that the average U-N distance is shorter than expected from a purely ionic bonding model. ... 

This is consistent with the diminished (in comparison to benzene) s-character of the C2-H bond in 1, as expressed by the large C-H coupling constant. The data are in agreement with a bonding model for cycloproparenes which uses sp-hybrids at the bridgehead positions and sp hybridization at C1. The chemical shifts and couplings of the other centers of the cycloproparenes are, however, almost unaffected. 

Results from local density models and BP, BLYP and EDF 1 density functional models are, broadly speaking, comparable to those from 6-3IG models, consistent with similarity in mean absolute errors. As with bond length comparisons, BLYP models stand out as inferior to the other non-local models. Both B3LYP/6-31G and MP2/6-31G models provide superior results, and either would appear to be a suitable choice where improved quality is required. 

Consistent with earlier remarks made for bond length comparisons, little if any improvement results in moving from the 6-3IG to the 6-311+G basis set for Hartree-Fock, local density and density functional models, but significant improvement results for MP2 models. 

The first section summarizes simply the essential features of the different types of molecular orbital calculations currently used to solve theoretical problems in organometallic chemistry. A critical comparison of these calculations is also given. The second section discusses the bonding in organometallic complexes and draws on recent computational results and develops the chemical and structural implications of bonding models based on perturbation theory arguments. 

Such a calculation had to be made in determining the problem of interaction between a diatomic polar molecule and an atom with a closed electronic shell [1,2]. The solution of this problem was of special importance in connection with the nature of the hydrogen bond. Our investigations were based on a definite hydrogen bond model that was verified by mathematical treatment and comparison with experimental results. This model is essentially different from that criticized by Dr. Burawoy. Its main point can be explained as follows [2, 3]. 

In summary, it is noted that multiple bonding between the heavier Group 14 elements E (Ge, Sn, Pb) differs in nature in comparison with the conventional a and 7T covalent bonds in alkenes and alkynes. In an E=Ebond, both components are of the donor-acceptor type, and a formal E=E bond involves two donor-acceptor components plus a p-p n bond. There is also the complication that the bond order may be lowered when each E atom bears an unpaired electron or a lone pair. The simple bonding models provide a reasonable rationale for the marked difference in molecular geometries, as well as the gradation of bond properties in formally single, double and triple bonds, in compounds of carbon versus those of its heavier congeners. 

Alkenes contain a C=C double bond. The C=C double bond can be described with two different models. According to the most commonly used model, a C=C double bond consists of a <7- and a tr-bond. The bond energy of the a-bond is 83 kcal/mol, about 20 kcal/mol higher than the tr-bond (63 kcal/mol). The higher stability of <7 bonds in comparison to n bonds is due to the difference in the overlap between the atomic orbitals (AOs) that form these bonds. Sigma bonds are produced by the overlap of two spn atomic orbitals (n 1,2,3), which is quite effective because it is frontal. Pi bonds are based on the overlap of 2,pz atomic orbitals, which is not as good because it is lateral or parallel. 

Minimization of electrostatic interactions as defined by the various parameters described above, was done in order to calculate bond angles in halomethanes, -silanes and -germanes, as well as one-centre compounds that contain lone pairs and double bonds. The comparison with experimental values is such as to leave no doubt about the power of the VSEPR model. 

A comparison of experimental values for intermetallic diatomic molecules with gold with the corresponding value calculated by the Pauling model and by the atomic cell model has been given in Table 6 of Reference ( ). Table 7 of Reference ( ) shows a comparison between experimental dissociation energies with values calculated by the atomic cell model and the empirical valence bond model. Table 9 of Reference ( ) takes Mledema s refinements (43) of the atomic cell model into account In these comparisons. 

These spectroscopic results are in agreement with the quantum mechanical calculations. CpaSi is used as a model for permethylsilicocene in order to reduce the number of internal coordinates It reacts as bis(monohapto)silicocene [5] with CO to a similiar adduct as found for the dimethylsilylene case The SiCO angle is calculated to be 159° and the silicon-carbon bond to be 210 pm, which underlines the reduced back-bonding in comparison to Me2SiCO The reaction enthalpy is, with -6 4 kcal/mole, nearly identical... 

Recently use of localised Wannier functions instead of delocalised Bloch states (CP orbitals) in CPMD simulations has proved an efficient and effective approach for study of fluids such as water133 and DMSO/water mixture.134 Use of localised Wannier functions also has the advantage that they allow electrons to be assigned to bonds, making visualisation of bonding and structure of molecules easy and facilitating comparisons with standard chemical bonding models. 

Accordingly, we identify 12.4 eV as the value of Eg from our bonding model. This closely agrees with Eg = 12.2 eV, found from the dielectric constant. Table 5.6 contains comparisons of the same kind for ionic compounds. The agreement is surprisingly good, since the excited states contributing to the polarizability need not be the same as the lowest excited states in the UV spectrum. 

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