Lecture #20, October 18, 2006

Cruising in to Chapter 16

Lecture Outline (some material shown is look-and-listen only) Quantum Theory of the Chemical Bond

Molecular orbitals (in "homonuclear diatomic molecules"

Approximated by combinations of atomic orbitals

Constructive and destructive interference effects

Combining 1s atomic orbitals

"Sigma" bonding molecular orbital

"Sigma" antibonding molecular orbital

If destructive interference represents the orbital in which the electron happens to be found, the linear combination results in a depeletion of electron density between the positively charged protons. They repel each other under these circumstances.  
A schematic of the orbitals on separated protons relative to "close" protons. The constructive combination of atomic orbitals gives rise to a system lower in energy than the isolated system and is a bonding molecular orbital.  
Energy diagram and geometry for the sigma-1s (bonding) orbital construction.
Re-defining how one calculates bond order within the context of molecular orbital theory.
The molecular orbital energy diagram for H2- and other isoelectronic species, all having bond order = 0.5. The electron configuration would be written as s1s2 s*1s2 s2s2
or, as in the current text, simply ss2 s*s2 ss2.
Energy diagram and geometry for the sigma*-1s (antibonding) orbital construction
Molecular orbital energy diagrams for 1, 2, 3, and 4 electrons systems. The correspondence between bond order and bondlength and between bond order and bond energy is shown for each molecular species.
The molecular orbital energy diagram for Li2 containing all six electrons. Since antibonding electrons negate the effect of bonding electrons, the pair of s* electrons cancels out the bonding characteristics of the pair of s electrons and both pairs revert to the inner core 1s2 configurations that we expect from simpler considerations. The electron configuration could thus be written as [He][He]s2s2.
The molecular orbital energy diagram for Be2. Since the number of antibonding electrons is equal to the number of bonding electrons, there is no net bonding in this molecule and it simply breaks back up into two Be atoms.
Energy diagram and geometry for the sigma-2px (bonding) orbital construction from the 2p atomic orbitals originally directed along the bonding (x) axis.
This (with luck) is an animation of the computer-generated s2p (or sp) bonding molecular orbital showing how the symmetry about the bond axis resembles that for the s-atomic orbital's symmetry about the nucleus. Hence the s label.
Energy diagram and geometry for the sigma*-2p (antibonding) molecular orbital construction
A reminder that, perpendicular to the bond axis in the x-direction, there are 2p atomic orbitals in the z-direction and in the y-direction. Just the 2py orbitals on each atom are shown here.
Energy diagram and geometry for the pi-2py (bonding) molecular orbital construction from the 2p atomic orbitals originally pointing in the y-direction, perpendicular to the bond axis and in the plane of the illustration. Constructive interference occurs in the region of overlap. (Recognize that there is an identical pi bonding level, same energy, constructed in the x-direction from the 2pz atomic orbitals.)
For "diboron", B2 we get a somewhat surprising result out of the molecular orbital energy diagram and the ensuing electron configuration. The ninth and tenth electrons go into the lowest energy available orbitals. These are the p orbitals. There are two of equal energy and Hund's rule forces us to put one electron in each and with parallel spins. Thus we have bond order = 1, a single bond, but it is not a sigma bond. In fact, there is one electron in one pi-orbital and another electron in a second pi-orbital comprising this single bond.