Ethylene orbital structure

Figure 1.14 The structure of ethylene. Orbital overlap of two sp hybridized carbons forms a carbon-carbon double bond. One part of the double bond results from a (head-on) overlap of sp2 orbitals (green), and the other part results from (sideways) overlap of unhybridized p orbitals (red/blue). The ir bond has regions of electron density on either side of a line drawn between nuclei.

In Ag-SAPO-ll/C2H4 zeolite the EPR at 77 K shows the spectra of Ag° atoms and C2H5 radicals. After annealing at 230 K those species disappeared and then an anisotropic EPR sextet was recorded. Based on DFT calculation the structure of complex was proposed in which two C2H4 ligands adopted eclipsed confirmation on either side of the Ag atom. As a result the overwhelming spin density is localised on ethylene orbitals. 

Fig. 2.2. Several representations of ethylene, which has two bonded carbon atoms, (a) is a ball and stick model, (b) shows a chemical structure, and (c) is a schematic orbital diagram. Note the trigonal planar arrangement of the bonds between the atoms-this arrangement is the consequence of the hybrid orbital structure.

ML2 complexes usually adopt a linear (or essentially linear) geometry (see 2.7). Knowledge of the orbital structure for strongly bent ML2 species is therefore useful mainly when one wishes to consider them as fragments in larger complexes. For example, the ethylene)Ni(PR3)2] complex (2-94) can be described in terms of a bent ML2 entity interacting with a molecule of ethylene. 

In the work we studied some features of formaldehyde oligomers and alkenes interaction resulting in the formation of 1,3-dioxanes. The geometric and orbital structures of the reactions transition states depending on the alkene structure are considered. The following alkenes were used ethylene (1), propylene (2), butene-1 (3), isobutylene (4) and t-butene-2 (5). A formaldehyde dimer (FD) served an example for FO calculation. 

HMO theory is named after its developer, Erich Huckel (1896-1980), who published his theory in 1930 [9] partly in order to explain the unusual stability of benzene and other aromatic compounds. Given that digital computers had not yet been invented and that all Hiickel s calculations had to be done by hand, HMO theory necessarily includes many approximations. The first is that only the jr-molecular orbitals of the molecule are considered. This implies that the entire molecular structure is planar (because then a plane of symmetry separates the r-orbitals, which are antisymmetric with respect to this plane, from all others). It also means that only one atomic orbital must be considered for each atom in the r-system (the p-orbital that is antisymmetric with respect to the plane of the molecule) and none at all for atoms (such as hydrogen) that are not involved in the r-system. Huckel then used the technique known as linear combination of atomic orbitals (LCAO) to build these atomic orbitals up into molecular orbitals. This is illustrated in Figure 7-18 for ethylene. 

The structure of ethylene and the orbital hybridization model for its double bond were presented m Section 2 20 and are briefly reviewed m Figure 5 1 Ethylene is planar each carbon is sp hybridized and the double bond is considered to have a a component and a TT component The ct component arises from overlap of sp hybrid orbitals along a line connecting the two carbons the tt component via a side by side overlap of two p orbitals Regions of high electron density attributed to the tt electrons appear above and below the plane of the molecule and are clearly evident m the electrostatic potential map Most of the reactions of ethylene and other alkenes involve these electrons... 

Structure. The straiued configuration of ethylene oxide has been a subject for bonding and molecular orbital studies. Valence bond and early molecular orbital studies have been reviewed (28). Intermediate neglect of differential overlap (INDO) and localized molecular orbital (LMO) calculations have also been performed (29—31). The LMO bond density maps show that the bond density is strongly polarized toward the oxygen atom (30). Maximum bond density hes outside of the CCO triangle, as suggested by the bent bonds of valence—bond theory (32). The H-nmr spectmm of ethylene oxide is consistent with these calculations (33). 

Many of the Lewis structures in Chapter 9 and elsewhere in this book represent molecules that contain double bonds and triple bonds. From simple molecules such as ethylene and acetylene to complex biochemical compounds such as chlorophyll and plastoquinone, multiple bonds are abundant in chemistry. Double bonds and triple bonds can be described by extending the orbital overlap model of bonding. We begin with ethylene, a simple hydrocarbon with the formula C2 H4. 

Now let us return to our discussion of the conical intersection structure for the [2+2] photochemical cycloaddition of two ethylenes and photochemical di-Jt-methane rearrangement. They are both similar to the 4 orbital 4 electron model just discussed, except that we have p and p overlaps rather than Is orbital overlaps. In Figure 9.5 it is clear that the conical intersection geometry is associated with T = 0 in Eq. 9.2b. Thus (inspecting Figure 9.5) we can deduce that... 

Consider first the ethylene molecule. Its geometrical structure is shown in Fig. 5. The s, py and pz atomic orbitals of the carbon atoms are assumed to be hybridized. This sp2 hybridization implies H-C-H bond angles of 120°, approximately in agreement with experimental results. The remaining two px orbitals are thus available to contribute to a -electron system in the molecule. Here again, the two linear combinations of atomic orbitals yield bonding and... 

As indicated above, the conventional structure of trans butadiene does not include the delocalization of the -electron system. This effect can be analyzed, at least approximately, by application of the Hfickel method. As in the example of ethylene, each carbon atom has an available p orbital-the... 

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