Molecular orbital theory: conjugation and aromaticity - Chemistry LibreTexts
The difference in benzene is that each carbon atom is joined to two other similar The six delocalised electrons go into three molecular orbitals - two in each. It uses 3-D pictorial presentations of molecular orbitals to elucidate organic reaction Lewis Theory uses curly arrows to denote electron migration during a chemical . Another important difference between Hydrogen Fluoride and previous . The upper bonding degenerate pair of orbitals are the HOMO's of benzene. Taking the difference of the two atomic orbitals, however, results in . In the VB description of the benzene molecule, each double bond is localized between a.
2.2: Molecular orbital theory: conjugation and aromaticity
Talk about a fucked-up building 2. Draw all the p orbitals with alternating phases. No two adjacent p orbitals should have lobes with the same phase. This is where things get interesting. In the case of cyclic systems, the n—1 rule fails. A meelion dollars if you can prove me wrong: We can, however, easily draw an orbital where all the phases alternate.
They are therefore the highest energy. For hexatriene, the second floor one node is fairly straightfoward: This is really the key difference in the molecular orbital picture of a cyclic system versus an acyclic system: For benzene, that results in a lowering of energy.31.02 Molecular Orbitals and Resonance Structures of Benzene
Any electron that occupies this orbital is excluded from the internuclear region, and its energy is higher than it would be if it occupied either atomic orbital. The molecular orbital energy-level diagramwhich is a diagram that shows the relative energies of molecular orbitals, for the H2 molecule is shown in Figure On either side of the central ladder are shown the energies of the 1s orbitals of atoms A and B, and the central two-rung ladder shows the energies of the bonding and antibonding combinations.
Only at this stage, after setting up the energy-level diagram, are the electrons introduced. In accord with the Pauli exclusion principleat most two electrons can occupy any one orbital. In H2 there are two electrons, and, following the building-up principlethey enter and fill the lower-energy bonding combination. Its low energy results in turn in the conventional interpretation, at least from the accumulation of electron density in the internuclear region because of constructive interference between the contributing atomic orbitals.
The central importance of the electron pair for bonding arises naturally in MO theory via the Pauli exclusion principle. A single electron pair is the maximum number that can occupy a bonding orbital and hence give the greatest lowering of energy. The molecular orbital energy-level diagram shown in Figure 13 also applies with changes of detail in the energies of the molecular orbitals to the hypothetical species He2.
The Pi Molecular Orbitals of Benzene
Although there is a bonding influence from the two bonding electrons, there is an antibonding influence from two antibonding electrons. As a result, the He2 molecule does not have a lower energy than two widely separated helium atoms and hence has no tendency to form. The overall effect is in fact slightly antibonding.
The role of the noble gas configuration now can be seen from a different perspective: The molecular orbitals of other species are constructed in an analogous way. In general, the orbitals in the valence shells of each atom are considered first not, initially, the electrons those orbitals contain. Then the sets of these orbitals that have the same symmetry with respect to the internuclear axis are selected.
This point is illustrated below. Bonding and antibonding combinations of each set are then formed, and from n atomic orbitals n such molecular orbitals are formed. The molecular orbital energy- level diagram that results is constructed by putting the molecular orbitals in order of increasing number of internuclear nodal planes, the orbital with no such nodal plane lying at lowest energy and the orbital with nodal planes between all the atoms lying at highest energy.
Molecular orbitals of period-2 diatomic molecules As a first illustration of this procedure, consider the structures of the diatomic molecules formed by the period-2 elements such as N2 and O2. Each valence shell has one 2s and three 2p orbitals, and so there are eight atomic orbitals in all and hence eight molecular orbitals that can be formed.
The energies of these atomic orbitals are shown on either side of the molecular orbital energy-level diagram in Figure It may be recalled from the discussion of atoms that the 2p orbitals have higher energy than the 2s orbitals. These four molecular orbitals lie typically at the energies shown in the middle of Figure The 2px orbitals on each atom do not have cylindrical symmetry around the internuclear axis.
The molecular orbital energy-level diagram for diatomic molecules of period-2 elements. The occupation of the orbitals is characteristic of N2. Now consider the structure of N 2. Note that only the orbitals in the lower portion of the diagram of Figure 14 are occupied.
Introduction to Molecular Orbital Theory
This configuration accounts for the considerable strength of the bonding in N2 and consequently its ability to act as a diluent for the oxygen in the atmospherebecause the O2 molecules are much more likely to react than the N2 molecules upon collision with other molecules.
To see how the MO approach transcends the Lewis approach and, in this instance, the VB approach as wellconsider the electronic configuration of O2. The other molecular orbitals are almost never drawn. The six carbon atoms form a perfectly regular hexagon. All of the carbon-carbon bonds have exactly the same lengths - somewhere between single and double bonds. There are delocalized electrons above and below the plane of the ring, which makes benzene particularly stable.
Because of the aromaticity of benzene, the resulting molecule is planar in shape with each C-C bond being 1. You might ask yourselves how it's possible to have all of the bonds to be the same length if the ring is conjugated with both single 1.
Interactions between Benzene Molecular Orbitals and Metal d Orbitals
Rather, the delocalization of the ring makes each count as one and a half bonds between the carbons which makes sense because experimentally we find that the actual bond length is somewhere in between a single and double bond. Finally, there are a total of six p-orbital electrons that form the stabilizing electron clouds above and below the aromatic ring. This further confirms the previous indication that the six-carbon benzene core is unusually stable to chemical modification.
The conceptual contradiction presented by a high degree of unsaturation low H: C ratio and high chemical stability for benzene and related compounds remained an unsolved puzzle for many years. Eventually, the presently accepted structure of a regular-hexagonal, planar ring of carbons was adopted, and the exceptional thermodynamic and chemical stability of this system was attributed to resonance stabilization of a conjugated cyclic triene. The High Stability of Benzene Here, two structurally and energetically equivalent electronic structures for a stable compound are written, but no single structure provides an accurate or even an adequate representation of the true molecule.