8.1 Valence Bond Theory

Chapter 8 contains advanced theories of bonding and orbitals to understand the observations.

A scientific theory is a strong explanation for observed natural laws or large bodies of experimental data.

It is necessary for a theory to explain experimental data and be able to predict behavior for it to be accepted.
The theory predicts three-dimensional molecular shapes that are consistent with experimental data collected for thousands of different molecules has gained widespread acceptance.
There is no explanation of chemical bonding provided by the VSEPR theory.

When atoms bond to form a molecule, atomic orbitals are not enough to describe the regions where electrons will be located.

A model that can account for the electronic structure of the molecule is required for a more complete understanding of electron distributions.
A popular theory holds that when a pair of electrons are shared by two atoms, they are attracted to each other by the nucleus of both atoms.
We will discuss how bonds are described in the following sections.

When two conditions are met, a covalent bond results when one atom overlaps the other, and the single electrons in each atom combine to form an electron pair.

The attraction between this negatively charged electron pair and the two atoms' positively charged nuclei serves to physically link the two atoms through a force we define as a covalent bond.
The strength of a bond depends on the amount of overlap.
Orbitals that overlap extensively form bonds that are stronger than those that have less overlap.

The sum of the energies is set by convention when the atoms are far apart.
The atoms begin to overlap as they move together.
The nucleus in the other atom attracts the electrons.
The nuclei and electrons begin to repel each other.
The attractions are stronger than the repulsions, and the energy of the system is decreasing.
The attraction of the nuclei for the electrons continues to increase as the atoms move closer together.
The energy reaches its lowest value at a specific distance between the atoms.
The bond distance between the two atoms is the optimum distance.
The bond is stable because the repulsive and attractive forces combine to create the lowest possible energy configuration.
The repulsions between the nuclei and the repulsions as electrons are confined in closer proximity to each other would become stronger if the distance between the nuclei were to decrease further.

The bond energy is the difference between the energy minimum and the energy of the two separated atoms.

When the bond is formed, this is the amount of energy released.
The same amount of energy is needed to break the bond.
The H2 molecule shown at the bond distance is 7.24 x 10-19 J lower in energy than the two separated hydrogen atoms.
This may seem small.
Bond energies are often discussed on a per-mole basis according to our earlier description of thermochemistry.
It takes 4.36 x 105 J to break 1 mole of H-H bonds, but it takes 7.24 x 10-19 J to break one H-H bond.
A comparison of bond lengths and energies is shown in We can find many of these bonds in a variety of molecule, and this table provides average values.
The first C-H bond in CH4 requires 439.3 kJ/mol, while the first C-H bond in H-CH2C6H5 requires 375.5 kJ/mol.

There is a chance of greater overlap if the orbitals are oriented in such a way that they overlap on a direct line between the two nuclei.

The structures suggest that O2 and N2 have double and triple bonds.
The triple bond consists of one s bond and two p bonds.
The first bond formed between any two atoms will always be a s bond, but there can only be one in any one location.
There will be one s bond and one or two p bonds in any multiple bond.
The bonds are described later in the chapter.

The average carbon-carbon single bond is 347 kJ/mol, while in a carbon-carbon double bond the p bond increases the bond strength by 267 kJ/mol.

A further increase of 225 kJ/mol is caused by adding an additional p bond.
When we compare other bonds, we can see a similar pattern.
Each individual p bond is weaker than a corresponding s bond between the same two atoms.
There is more overlap in a s bond than in a p bond.