Orbital Theories

like ionic compound with a high melting point, while methane, being non-salt-like, has a boiling point of -161 degrees Celsius. Methane or CH4 is mainly a covalently-bonded molecule. There is little electrostatic attraction between methane molecules (they are nonpolar) so that they boil at a very low temperature to become a gas.

Hydrogen fluoride boils, on the other hand, 200 degrees higher than methane gas. This is because it is more ionic and has connections between the Hydrogen and Fluorine atoms through not only ionic bonding but also with hydrogen bonding between the HF molecules in chains and rings. This makes the boiling point higher.

Because methane is nonpolar, it is inert to almost all reagents that could remove the hydrogen ion under anything but the most extreme conditions. It is therefore difficult to generate a CH3+ or CH3- ion, as these would be very reactive and just wouldn’t last long. In other words, you couldn’t make CH4 plus HF to make CH3F plus H2. This would require a cation of CH3+, which is not very stable in nature.

ORBITAL THEORIES

While there are many ways to describe and write covalent bonds (which will be described later in this chapter), there are some concrete theories about how covalent bonds exist in organic molecules. While there is ionic bonding in organic chemistry, the most important bonding in this type of chemistry involves covalent bonds. You may have learned that covalent bonding involves the sharing of electrons but you may not know exactly how this happens.

This introduces the topic of the valence bond theory, which describes how bonding happens in covalent molecules. According to this theory, there are two atomic orbitals around a pair of atoms—each orbital of which contains one electron. In sharing orbitals, the orbital pair will contain a stable set of two electrons.

The H2 molecule is the simplest case of a covalently-bonded molecule. There are two spherical orbitals (1s orbitals) in each hydrogen atom, each with one electron in it. In the bonding of the two atoms, the electrons no longer are anywhere within the sphere but spend more time in that part of each sphere between the two nuclei, which holds the

atoms together. According to quantum chemistry, the two electrons must occupy a shared orbital space in order to form a bond.

The hydrogen atoms are too far apart, the 1s orbitals cannot overlap in order to form a covalent bond. As they overlap, a bond will begin to form, lowering the potential energy of the system. There will be attraction of the opposite electron to the opposite nucleus. They cannot, however, get too close or there will be repulsion of the nuclei with each other, which increases the potential energy of the system. A balance is had when there is an optimal distance between the nuclei.

In a sense, there is a defined “optimal” distance between two bonding nuclei, which is when the potential energy is the lowest, meaning that the combined repulsive and attractive forces add up to the greatest amount of attraction. This optimal internuclear distance is referred to as the “bond length”. The difference between the optimal “bonded” energy and the separated energy is the “bond energy”.

Every covalent bond has a certain strength and bond length. As an example, the carboncarbon bond is 1.5 Angstroms long with an Angstrom being 10-10 meters in length. Bonds aren’t the straight sticks as they are depicted in drawing chemical bonds. Instead, they are more like springs, which have the ability to bend, extend, and compress.

According to the valence bond theory, there are two assumptions about bonds: 1) the strength of a given bond is directly proportional to the degree of overlap of the bond (in other words, the greater the overlap, the more stable is the bond). 2) An atom is able to use different combinations of atomic orbitals in order to maximize the overlap of the orbitals used by bonded atoms. Maximum overlap occurs between orbitals of the same spatial orientation and similar energies.

Bonding can take place between beryllium and hydrogen to make BeH2 which, according to the periodic table of the elements in figure 3, has an atomic number of 4:

Figure 3

According to the atomic number and the atomic orbital approach, beryllium has a 1s22s2 electron configuration. It has apparently filled its 2s orbital subshell, leaving behind no apparent reason why it should want to overlap with the singly occupied 1s orbital associated with the hydrogen atom. How do these two atoms overlap to make BeH2? One way to do this is to excite a 2s electron on beryllium to allow for a partially empty 2p orbital that absolutely could bind with hydrogen. Doing this is called “promotion”. What this looks like is seen in figure 4:

Figure 4.