What are Molecular Orbitals? It is clear that ions are held together by electrostatic attraction in an electrovalent or ionic bond, but it is less easy to understand why a shared pair of electrons should hold two atoms together. There are two ways of viewing the problem. One involves the use of the concept of resonance; the other line of approach is afforded by the method of molecular orbitals. In this method, a molecule is supposed to have electronic orbitals associated with it in much the same way as a single atom has. All, or nearly all, of the electrons associated with the component atoms of a molecule are supposed to enter the various molecular orbitals, which fill up according to certain rules just as in atomic orbitals and no molecular orbital can contain two precisely similar electrons. This means that any particular molecular orbital can only contain two electrons, and these two must have different spins.
The nomenclature s, p and d used for atomic orbitals is replaced by that of σ, π and δ for molecular orbitals, σ and π being the most important. Electrons in some of these molecular orbitals contribute towards the binding together of the atoms in a molecule; such orbitals are known as bonding orbitals. Others cause repulsion between atoms and are known as anti-bonding orbitals.
The nomenclature most commonly used indicates (a) the nature of the molecular orbital, i.e. σ or π, occupied by electrons in a molecule, (b) the atomic orbital from which the electrons originated, and (c) whether the molecular orbital is bonding or anti-bonding. A σIs orbital, for example, is a σ orbital made up by interaction of Is atomic orbital’s; it is a bonding orbital. A σ* Is orbital is the corresponding anti-bonding orbital.
The hydrogen molecule. In a hydrogen molecule, H2 the two I S electrons, one from each of the two hydrogen atoms concerned, are present in a σ Is molecule orbital. This is a bonding orbital and constitutes the single covalent bond in the hydrogen molecule. Since the orbital can only hold two electrons if they have different spins, it is clear that a molecule will only be formed from two hydrogen atoms containing electrons with opposed spins. This is an important point. A covalent bond is not simply a shared pair; it is a shared pair of electrons with opposed spins. This means that only single, unpaired electrons, in atoms, can participate in covalent bond formation.
An atomic orbital can be represented as a charge cloud of varying density, or more conveniently, by mapping out the boundary surface within which the electron might be said to exist. The same procedure can be adopted for molecular orbitals, as is illustrated for the hydrogen molecule.
It is the accumulation of negative charge between the two positively charged atomic nuclei which is responsible for holding them together. Such an accumulation of charge occurs when two s orbitals containing electrons with opposed spins overlap.
Two Is atomic orbitals; A σ Is molecular orbital.
Electrons with different A bonding orbital forming a σ-bond
Formation of a σ-bond from a pair of s electrons with different spins.
If electrons with the same spin are involved, an anti-bonding molecular orbital is formed.
The bonding molecular orbital formed from any pair of s electrons with different spins is plum-shaped, as shown in the hydrogen molecule. It is symmetrical about the line joining the two nuclei and has no nodal plane,
Two IS atomic orbitals; A σ* Is molecular orbital.
Electrons with like spins Anti-bonding.
(Formation of anti-bonding molecular orbital from a pair of s electrons with like spins).
i.e. a plane in which the probability of finding an electron is zero. The bond which it forms is referred to as a σ-electrons.
Two p atomic orbitals. A π molecular orbital
Formation of a π molecular orbitals from a pair of p electrons with different spins. A π-bond.
π-orbitals, like p atomic orbitals from which they are formed, have a nodal plane. The bonds which they form are known as π-type or π bonds and electrons occupying a – orbital are known as π-electrons.