Qualitative Molecular Orbital Theory

Molecular orbital theory (MO theory or MOT) is a way of employing quantum mechanics to describe the electronic structure of molecules in chemistry. Early in the twentieth century, it was proposed.

Electrons in a molecule are not ascribed to distinct chemical interactions between atoms in molecular orbital theory, but are instead viewed as circulating under the influence of the atomic nuclei across the molecule.

As molecular orbitals which encompass two or more atoms in a molecule and include valence electrons between atoms, quantum mechanics describes the spatial and kinetic properties of electrons.

By estimating the states of bonded electrons—the molecular orbitals—as linear combinations of atomic orbitals, molecular orbital theory revolutionised chemical bonding research (LCAO).

Density functional theory (DFT) or Hartree–Fock (HF) models are used to approximate the Schrödinger equation.

The molecular orbitals can be obtained using the “Linear combination of atomic orbitals molecular orbital method” for an imperfect, yet qualitatively relevant, explanation of the molecular structure. The molecular orbitals are written as linear combinations of atomic orbitals in this case.

Developed MOT

Friedrich Hund, Robert Mulliken, John C. Slater, and John Lennard-Jones contributed to the development of molecular orbital theory in the years following valence bond theory was established (1927). Originally, the HundMulliken hypothesis was known as the MO theory.

The theory of quantum mechanics was created by Erwin Schrödinger in 1926.

Mulliken and Friedrich Hund collaborated on a quantum interpretation of diatomic molecules’ spectra.

Their molecular orbital hypothesis, which involved assigning electrons to states that stretched throughout a whole molecule, was published in 1927.

In 1931, Hund was the first to mention bonds, and in 1932, Mulliken was the first to use the term orbital. The Hund-Mulliken hypothesis was recognised as a legitimate and useful theory by 1933.

The main points of molecular orbital theory

1. Atomic orbitals include electrons, while molecular orbitals contain electrons.
2. The molecular orbitals are made up of atomic orbitals with similar energies and symmetry.
3. The BMO is more stable than the comparable ABMO because it has less energy.
4. The Aufbau principle, Pauli’s exclusion principle, and Hund’s rule are all used to fill the molecular orbitals with electrons.
5. Linear combinations of atomic orbitals with almost equal energies constitute molecular orbitals.
6. The nuclei of bound atoms in a molecule are related with Molecular orbitals.
7. The number of merging atomic orbitals equals the number of molecular orbitals that are created.
8. The freshly generated molecular orbitals are where a molecule’s electrons are accommodated. In the same way that electrons are filled in atomic orbitals, the filling of these orbitals follows Aufbau’s principle, Pauli’s exclusion principle, and Hund’s rule.
9. The number of covalent bonds formed by the two merging atoms is determined by the bond order. The following equation can be used to figure out a molecule’s bond order.
10. The number of combining atomic orbitals equals the number of molecular orbitals created. The energy of half of the molecular orbitals created will be lower than the corresponding atomic orbital, while the energy of the remaining molecular orbitals will be higher.
11. The bonding molecular orbital has the lowest energy, while the anti-bonding molecular orbital has the highest. (Sigma), (pi), and (delta) are the bonding molecular orbitals, while σ*, π* and δ* are the antibonding orbitals.

The molecular orbitals of molecular theory

To comprehend the bonding of a diatomic molecule, use a diatomic molecular orbital diagram. MO diagrams can be used to figure out a molecule’s magnetic characteristics and how they change as the molecule is ionised. They also reveal the molecule’s bond order and how many bonds the two atoms share.

By applying the Schrödinger equation to a molecule, the energy of the electrons can be better understood. Quantum mechanics can precisely describe the energies of single electron systems, but the Born-Oppenheimer Approximation may be used to approximate the energies of multiple electron systems, assuming the nuclei remain stationary. To further describe the state of the molecule, the LCAO-MO technique is employed in conjunction.

Only two atoms are linked together in a diatomic molecule. Homonuclear and heteronuclear nuclei are the two types of nuclei. A homonuclear diatomic molecule is made up of two identical atoms. H2, O2, and N2 are some examples of these gases. Two atoms from two distinct elements make up a heteronuclear diatomic molecule.

The and molecular orbitals are formed by superpositioning the two 1s atomic orbitals. The orbital is formed when two atomic orbitals are in phase, resulting in a higher electron density. A node between the two 1s orbitals creates a jump in energy, the * orbital, if they are not in phase.

The bond arrangement and number of bonds formed between the two atoms can be deduced from the diagram. It is equal to one for this molecule. If a molecule is ionised, bond order can also reveal how near or stretched a bond is. 

Conclusion

The atomic orbitals are merged to generate molecular orbitals, according to molecular orbital theory. The energy of the electrons is reduced because the electron density of each atom is spread out over the entire molecule. This explains why bonding results in stabilisation. The level of stabilisation is determined by the amount of overlap and energy difference between atomic orbitals. Stable molecular orbitals are formed when atomic orbitals overlap. The overlapping atomic orbitals must have equal energies as one of the conditions for overlap.

To demonstrate how molecular orbital theory works, we’ll look at how homonuclear and heteronuclear diatomic molecules bind.

Because characterising molecular orbitals in polynuclear compounds is difficult, we shall use the concept of bonding through hybridised atomic orbitals to account for bonding in such systems.

The molecular orbital theory can tell us about both ionic and covalent compounds, and it can also predict which ones will be ionic and which will be covalent. Chemists may use it to anticipate the properties of molecules, and it’s a powerful and complicated instrument.