Carbon atom orbital energies

Orbital correlation diagram for carbon monoxide. The carbon atomic orbital energies are on the left, and the oxygen atomic orbital energies are on the right. The molecular orbitals that form from mixing of the atomic orbitals are represented by the horizontal lines in the center at their approximate orbital energies in the CO molecule. The vertical lines indicate the orbital occupancy. 

Figure 21-13 Calculated molecular-orbital energies of planar cyclic systems made up of N 2p carbon atomic orbitals, in units of a and fi.

As may be seen in figure 4 and as noted above for C60 the fullerene n molecular orbitals are composed of carbon atomic orbitals which contain a substantial amount of 2s orbital character and as the carbon 2s orbital lies so much lower in energy than the 2p orbital the resulting orbitals are expected to exhibit an enhanced electronegativity when compared with their planar counterparts. Furthermore the fullerenes are not only of intermediate hybridization, but probably exhibit variable hybridization as they pass through different states of reduction, for the pyramid-alization of anions is well known. Based on figure 3 it seems that the reduction potential of C60 is lowered by about 0.8V as a result of rehybridization effects. 

The special stability of benzene (aromaticity) comes from the six tc electrons in three molecular orbitals made up by the overlap of the six atomic p orbitals on the carbon atoms. The energy levels of these orbitals are arranged so that there is exceptional stability in the molecule (a notional 140 kjmor1 over a molecule with three conjugated double bonds), and the shift of the six identical hydrogen atoms in the NMR spectrum (5h 7.2 p.p.m.) is evidence of a ring current in the delocalized tc system. 

The most probable excitations in the valine fi decay correspond to transitions—of an electron from the MO made up of the 2s carbon atom orbitals and the nearly orbitals of the hetero atoms, and of the Is orbitals of helium—into the lower vacant MO, and they have an energy of 40 eV. 

The a bonds in the backbone of vinyl polymers should not be describable in terms of local states of small model molecules because of overlap of carbon atomic orbitals only 1.5X apart. This concept can be tested in polyethylene where the least bound C-C bond band widths have been calculated to be about 3 eV (.8). The energy loss function, Im(-l/e), for polyethylene is given in Figure 3 where and the real and imaginary parts of the... 

The n bonding orbital is polarized toward the oxygen atom. The results of an SHMO calculation yield n = 0.540(2/ , ) + 0.841 [2pq). Thus the smaller coefficient of the carbon 2p orbital means that the carbonyl n bond will interact rather weakly with substituents attached to the carbon atom. The energy of the HOMO, = a — 1.651 jl, suggests that... 

Figure 10.8 Electronic structures of CO interacting with a Rul9 cluster. Overlap population densities of states are represented of the carbon atomic orbitals of CO interacting with Ru d-valence electrons as a function of electron orbital energies. Orbital densities are shown for energies of maximum electron...

The orbital combinations for p and Py (hereafter abbreviated as x and y, respectively), the p orbitals perpendicular to the chain propagation axis, are of course of the same energy (or degenerate). We see from 19 that s, x, and y give the most bonding crystal orbital when the unit cell (here a single carbon atom) orbitals are all in-phase, i.e., have the same sign. This is the k = 0 (termed r) Bloch function. 

The first RE is of limited interest, as many molecules cannot absorb an energy of the order of 54 eV in an ionization process. This is probably the case for all hydrocarbons, as they are built up of carbon atomic orbitals with orbital energies of the order of 300, 24, and 13 eV, and hydrogen atomic orbitals at 13 eV. As the third RE is too low for ionization of most molecules, the electron will enter the 2s or 2p orbital of He if partial charge exchange takes place when He meets a molecule. The effective RE will then be somewhat lower than 13.6 eV. 

Such geometry implies sp hybridization for each carbon atom. The energy levels of the atomic orbitals of the sp hybridized carbon atom are shown in Fig. 19.7c. The orientation of the sp and p orbitals is shown in Fig. 19.8a. 

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