Ethylene hybrid atomic orbitals

Similar, but different, redeployment is envisaged when a carbon atom combines with three other atoms, e.g. in ethene (ethylene) (p. 8) three sp2 hybrid atomic orbitals disposed at 120° to each other in the same plane (plane trigonal hybridisation) are then employed. Finally, when carbon combines with two other atoms, e.g. in ethyne (acetylene) (p. 9) two sp1 hybrid atomic orbitals disposed at 180° to each other (idigonal hybridisation) are employed. In each case the s orbital is always involved as it is the one of lowest energy level. 

A possible explanation for the calculated trend in the metal ion binding energies to ethylene, which will be discussed for the triply bonded substrates, could lie in a consideration of the Dewar-Chatt-Duncanson donor-acceptor model for bridging-type metal-olefin complexes. Their proposed two-way interaction involves mixing of the olefin n electrons with a metal (n + l)sp a hybrid atomic orbital (L —> M, for short) and simultaneous back donation (M L) of metal nd electrons of appropriate symmetry into the olefin k molecular orbital MO. For the monocation metal ions the latter-type interaction should be less favourable due to stabilizaion of the nd electrons by the charge on the metal. L M should be favoured for the same reason stabilization of the (n + l)s and (n+ l)p orbitals by the + 1 charge. 

To summari2B, we have seen that the carbon atoms of ethylene are connected via a a bond and a it bond. The a bond results from the overlap of r -hybridized atomic orbitals, while the It bond results from the overlap ofp orbitals. These two separate bonding interactions (n and it) comprise the double bond of ethylene. 

Ethylene is planar with bond angles close to 120° (Figure 2 15) therefore some hybridization state other than sp is required The hybridization scheme is determined by the number of atoms to which carbon is directly attached In sp hybridization four atoms are attached to carbon by ct bonds and so four equivalent sp hybrid orbitals are required In ethylene three atoms are attached to each carbon so three equivalent hybrid orbitals... 

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... 

Boron trifluoride and ethylene are but two of the many instances where the directional properties of covalent bonds are better described in terms of overlap of hybrid orbitals than in terms of simple atomic orbitals. 

All four carbon atoms of buta-1,3-diene are sp2 hybridized, and (in the planar conformation) they all have overlapping p orbitals. Let s review how we constructed the pi molecular orbitals (MOs) of ethylene from the p atomic orbitals of the two carbon atoms (Figurel5-3). Each p orbital consists of two lobes, with opposite phases of the wave function in the two lobes. The plus and minus signs used in drawing these orbitals indicate the phase of the wave function, not electrical charges. To minimize confusion, we will... 

The pi molecular orbitals of ethylene. The pi bonding orbital is formed by constmctive overlap of unhybridized p orbitals on the sp2 hybrid carbon atoms. Destructive overlap of thesep orbitals forms the antibonding pi orbital. Combination of two atomic orbitals must give exactly two molecular orbitals. 

The chemical bonding of re coordinated ethylene is very similar to the Chatt-Dewar-Duncanson picture of CO coordination (Fig. 4.6). The donating orbital is the doubly occupied n orbital that is a-symmetric with respect to the normal to the surface. When adsorbed atop it interacts with the highly occupied dz2 surface atomic orbital and the partially filled s and pz orbitals. The ethylene LUMO is the empty asymmetric n orbital, which interacts with the surface dxz and px orbitals. The corresponding overall interaction is relatively weak and no hybridization on ethylene is assumed to occur. The orbital interaction diagram of the occupied ethylene n orbital with surface atom dz2 orbital is analogous to that sketched for the CO 5a orbital in Fig. 4.4. When this dz2 becomes nearly completely occupied, as occurs, for instance, for Pd or Pt, the ethylene-rc surface atom dz2 interaction... 

Atomic orbitals that do not participate in the hybridization are then used for bonding with other atomic orbitals on adjacent centers as long as there is nonzero overlap of the atontic orbitals. For example, a hybridization description for the bonding in ethylene accounts for a so-called sigma (a) framework of bonding utihzing sp hybrids and a second interaction called a pi (tt) interaction between pure p atontic orbitals on the carbon atoms (8). 

In unsaturated compounds another type of hybridization makes its appearance, and is known as an sp hybrid. In this case the carbon atom retains one p atomic orbital containing a single electron which lies above and belowthe plane of the three hybrids, see (D) in Fig. 12.3. In ethylene, two rp orbitals are combined with hydrogen and one in each carbon atom forms a mutual link. This leaves the molecule with two 2p orbitals which unite to form a it orbital (see Fig. 12.4). 

The structure of ethylene and the orbital hybridization model for the double bond were presented in Section 1.17. To review. Figure 5.1 depicts the planar structure of ethylene, its bond distances, and its bond angles. Each of the carbon atoms is xp -hybridized, and the double bond possesses a o component and a tt component. The o component results when an sp orbital of one carbon, oriented so that its axis lies along the intemuclear axis, overlaps with a similarly disposed sp orbital of the other carbon. Each sp orbital contains one electron, and the resulting a bond contains two of the four electrons of the double bond. The tt bond contributes the other two electrons and is formed by a side-by-side overlap of singly occupied p orbitals of the two carbons. 

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