8.1 Valence Bond Theory

The VSEPR theory described in the previous chapter is widely accepted because it provides 3D shape predictions that align closely with experimental data. However, this theory does not address the details of chemical bonding. Two theories of bonding have emerged and understanding both is beneficial to understanding the formation of chemical bonds.

1. Valence Bond Theory - a description of a covalent bond as the overlap of two atomic . . . a specified volume of atomic space where an electron can be found 95% of the time. each containing one electron.

   In Chapter 7, we learned that the two hydrogen atoms in a hydrogen molecule, H₂, are held together with a covalent bond. The animation to the right depicts the energy involved in the . . . energy must be added to break chemical bonds (an endothermic process, +ΔH), whereas forming chemical bonds releases energy (an exothermic process, –ΔH).. For the two separated hydrogen atoms to bond, their internuclear distance must decrease.

    As the hydrogen atoms approach each other, their valence electrons (1s¹) begin to interact with each other and establish an attractive force with the nuclei of both atoms.   At an internuclear distance of 74 pm, the H–H bond has fully formed and both electrons freely move throughout the space around the H₂ molecule. The H–H bondlength of 74 pm corresponds to an energy minimum of –7.24 × 10^-19 J.   Decreasing the internuclear distance further causes a rise in the potential energy as the positive charges in the two nuclei begin to repel each other. 
   
   The bond length is defined as the internuclear distance at which the lowest potential energy is achieved (74 pm for H–H). The bond energy is the difference between the energy minimum (which occurs at the average bond distance . . . . 74 pm for H–H) and the energy of the two separated atoms. This is the quantity of energy released (exothermic, –ΔH) when the bond is formed and absorbed (endothermic, +ΔH) when the bond is broken.

   While the animation above describes much of the bonding process, it does not address where the bonding electrons are located most of the time. In the valence bond theory, this "new orbital" is the sum of the two original atomic orbitals. As the two atomic orbitals begin to overlap (combine), the individual electrons "pair" - one with a m_s of +½ and the other with a m_s of –½. This negatively-charged electron pair is simultaneously attracted to the positive nuclei of both contributing atoms, resulting in the formation of a stable chemical bond.

2. Molecular Orbital Theory - describes bonding through the creation of molecular orbitals. This theory asserts that atomic orbitals are present in atoms and molecular orbitals are present in molecules. As two atoms approach each other to initiate bond formation, a region in space that is similar to, but different from, the original atomic orbitals becomes "home" to the newly paired electrons. The first bond formed between two atoms is called a sigma bond, σ. The molecular orbital of a σ bond is localized, meaning that it is in the vicinity of, or local to, the two bonded atoms. The . . . a double (=) bond, and . . . a triple (≡) bond are called pi bonds, π. The molecular orbital of π bonds can be delocalized - the orbital can encompass more than two atoms and in some cases a molecular orbital may contain all the atoms of a molecule.

Both theories provide different, useful ways of describing molecular structure. Of course, the location and size of the molecule's orbitals are determined based on energy - a molecule's bonding electrons occupy spatial regions where the repulsion energy between bonds is the lowest. If an orbital structure where all the electrons are localized results in the lowest energy, then that is what occurs. However, for many molecules that contain double and triple bonds, the lowest energy is achieved when electrons travel in larger (delocalized) orbitals that keep the electron pairs (bonding and nonbonding) away from each other while maintaining bond integrity.

Sigma bonds are oriented along the line formed by the two bonded atoms while pi bonds are oriented perpendicular to sigma bonds. As a result, sigma bonds form from head-to-head overlap of orbitals while pi bonds form from sideways overlap of orbitals.