- As we know that Lewis approach helps in writing the structure of molecules but it fails to explain the formation of the chemical bond. It also does not give any reason for the difference in bond dissociation enthalpies and bond lengths in molecules like H2 (435.8 kJ mol-1, 74 pm) and F2 (155 kJ mol-1, 144 pm), although in both cases a single covalent bond is formed by the sharing of an electron pair between the respective atoms. It also gives no idea about the shapes of polyatomic molecules.
- Similarly, the VSEPR theory gives the geometry of simple molecules but theoretically, it does not explain them and also it has limited applications. To overcome these limitations the two important theories based on quantum mechanical principles are introduced. These are the valence bond (VB) theory and molecular orbital (MO) theory.
Valence Bond Theory
- Valence bond theory was introduced by Heitler and London (1927) and developed further by Pauling and others.
- A discussion of the valence bond theory is based on the knowledge of atomic orbitals, electronic configurations of elements (Units 2), the overlap criteria of atomic orbitals, the hybridization of atomic orbitals and the principles of variation and superposition.
To start with, let us consider the formation of a hydrogen molecule which is the simplest of all molecules:
- Consider two hydrogen atoms A and B approaching each other having nuclei NA and NB and electrons present in them are represented by eA and eB.
- When the two atoms are at a large distance from each other, there is no interaction between them.
- As these two atoms approach each other, new attractive and repulsive forces begin to operate.
- Attractive forces arise between:
(i) The nucleus of one atom and its own electron, i.e. NA – eA and NB– eB
(ii) The nucleus of one atom and electron of the other atom, i.e. NA– eB and NB– eA
- Similarly, repulsive forces arise between:
(i) electrons of two atoms like eA – eB
(ii) nuclei of two atoms NA – NB
- Attractive forces tend to bring the two atoms close to each other whereas repulsive forces tend to push them apart.
Forces of attraction and repulsion during the formation of H2 molecule
- Experimentally it has been found that the magnitude of new attractive force is more than the new repulsive forces.
- As a result, two atoms approach each other and potential energy decreases. Ultimately a stage is reached where the net force of attraction balances the force of repulsion and the system acquires minimum energy.
- At this stage, two hydrogen atoms are said to be bonded together to form a stable molecule having a bond length of 74 pm.
- Since the energy gets released when the bond is formed between two hydrogen atoms, the hydrogen molecule is more stable than that of isolated hydrogen atoms. The energy released is known as bond enthalpy, which is corresponding to the minimum in the curve depicted in the graph below.
- Conversely, 435.8 kJ of energy is required to dissociate one mole of H2 molecule.
H2(g) + 435.8 kJ mol–1 → H(g) + H(g)
Potential energy curve for the formation of H2 molecule as a function of the internuclear distance of the H atoms
Molecular Orbital Theory
- There is another approach to chemical bonding known as molecular orbital theory (MOT) developed by Mulliken (1932) and Hund, which explains the bonding characteristics in a better way.
- The molecular orbital theory considers the entire molecule as a unit with all the electrons moving under the influence of all the nuclei present in the molecular.
- This approach recognizes that each electron belongs to the molecule as a whole and may move within the entire molecule.
- When the atoms to be bonded come close together, the orbitals of the bonded atoms lose their individual character and fuse (overlap) to form larger orbitals called molecular orbitals.
- Like atomic orbitals, there are molecular orbitals in a molecule. The only difference is that in atomic orbitals, electrons move under the influence of only one nucleus (i.e. Atomic orbital are monocentric), while in molecular orbitals, electrons move under the influence of many nuclei, they are polycentric.
Important features of M.O.T.
(i) Like an Atomic orbital which is around the nucleus of an atom there is a molecular orbital which is around the nuclei of a molecule.
(ii) The molecular orbitals are entirely different from the atomic orbitals from which they are formed.
(iii) The valence electrons of the constituent atoms are considered to be moving under the influence of nuclei of participating atoms in the molecular orbital.
(iv) The molecular orbitals possess different energy levels like atomic orbitals in an isolated atom.
(v) The shape of molecular orbitals are dependent upon the shapes of atomic orbitals from which they are formed.
(vi) Molecular orbitals are arranged in order of increasing energy just like atomic orbitals. (vii) The number of molecular orbitals formed is equal to the number of atomic orbitals combining in bond formation.
(viii) Like atomic orbitals, the filling of electrons in molecular orbitals is governed by the three principles such as the Aufbau principle, Hund’s rule and Pauli’s exclusion principle.
Conditions for atomic orbitals to form M.O.
- The combining A.O. should be of comparable energy.
- The combining atomic orbitals must overlap to a large extent greater the overlap, stable is the molecule formed.
Relative energies of M.O. and filling of electron
Energy diagram is shown below:
M.O Energy level diagram for O2, F2 and Ne
M.O energy diagram for Li2, Be2, B2, C2 and N2 molecule