Diatomic Molecules with s and p Valence Atomic Orbitals
With the second row elements we have both s and p valence orbitals. These overlap to form molecular orbitals. The same approach that was used with 1s orbitals, the adding and subtracting of atomic orbitals, is used to describe bonding and antibonding molecular orbitals. The 2s orbitals are added, as with the H2 molecule, to form a lower energy sigma bonding molecular orbital. The difference between 2s orbitals gives the higher energy sigma antibonding orbital.
The 2pz atomic orbitals are oriented along the internuclear axis. They overlap to form a lower energy sigma bonding orbital with nodal planes at each nuclei and a sigma antibonding molecular orbital with nodal planes at the two nuclei and midway between the two nuclei. The remaining 2px and 2py atomic orbitals overlap on either side of the Z axis. These form two pi bonding and two pi antibonding molecular orbitals. When rotated about the Z axis, a pi orbital appears identical after rotating 180 degrees. The px molecular orbitals are at right angles to the py orbitals. An energy level diagram is shown below.
If there is a small sedation in the energy of the atomic 2s and 2p orbitals there will be an interaction between the sigma s bonding orbital and the pz atomic orbitals. This is called hybridization. This hybridization is shown in the diagram below. The sigma z bonding orbital is at a higher energy than the pi x and pi y bonding orbitals, the reverse of the previous diagram where the energies of the 2s and 2p atomic orbitals are at greater energy separation.
The shapes of these orbitals are shown in the following diagram. These shapes represent the surface that encloses 99% of the electron probability density. The pyb and py* orbitals are not shown. They have the same form as the pxb and px* orbitals but rotated 90 degrees about the z axis.
The lithium atom has one 2s valence electron. There is little difference in the energy of the 2s and 2p orbitals, thus there is considerable hybridization between the sigma s bonding molecular orbital and the sigma z bonding orbital in the Li2 molecule. The Li2 molecule has two electrons which form an electron pair in the sigma s bonding orbital. The bond energy is 26.3 kcal mole-1 and the bond length is 2.67 Angstroms. The bond length is greater than with H2 largely because of the larger radius of the 2s orbital than the 1s orbital of hydrogen.
Be does not form a diatomic molecule. This is because such a molecule would have two electrons in the sigma s bonding orbital and two electrons in the sigma antibonding orbital. The total number of bonding electrons would be zero.
Boron forms the B2 molecule with six electrons. The first four electrons are similar to the four electrons in the hypothetical Be molecule. The next two electrons occupy the pi x and pi y bonding orbitals of the hybridized energy level model.
The C2 molecule has the ground-state structure (ssb)2 (ss*)2 (pxb)2 (pyb)2. There are no unpaired electrons. There are two net bonds. The bond energy is 144 kcal mole-1 and the bond length is 1.24 Angstroms.
N2 has the ground-state configuration of (ssb)2 (ss*)2 (pxb)2 (pyb)2 (szb)2. This means a net bond order of three, the maximum for a second row diatomic molecule. This triple bond accounts for the well known stability of the N2 molecule.
O2 has the ground-state configuration (ssb)2 (ss*)2 (szb)2 (pxb)2 (pyb)2 (px*)1 (py*)1. By Hund’s rule the two p* electrons have the same spin in the ground state, thus there are two unpaired electrons. Accordingly O2 is paramagnetic. The bond energy is 118 kcal mole-1 and the bond length is 1.21 Angstroms.
The electronic configuration of F2 is (ssb)2 (ss*)2 (szb)2 (pxb)2 (pyb)2 (px*)2 (py*)2. This gives a net of one bond. The molecule is diamagnetic. The bond energy is 33 kcal mole-1 and the bond length is 1.42 Angstroms.
Neon doesn’t form an Ne2 molecule. The hypothetical molecule would have a net bond order of zero.
With some adjustment the molecular orbital energy level diagrams used for diatomic molecules of second row elements applies to larger atoms as well.