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Limitations of Lewis approach
It failed to explain:
Formation of a chemical bond
Shapes of polyatomic molecules
Difference between bond dissociation energies and bond lengths of various molecules containing single kind of bonds. For example, both H2 and F2 contain a single covalent bond, but have different values of the same quantities.
Limitations of VSEPR theory
Although it gave the geometry of simple molecules, it could not explain them.
It had limited applications.
To overcome these limitations, two new theories were given.
Valence bond theory
Molecular orbital theory
Valence Bond Theory
Introduced by Heitler and London (1927); developed further by Pauling and others
Consider two hydrogen atoms A and B approaching each other to form a covalent bond. NA, NB, eA, and eB represent the nucleus and electrons of atoms A and B respectively.
When the atoms are at large distance from each other, there is no force of attraction between them.
As the distance reduces, the following forces start developing in them:
Attractive forces −
1. Between the nucleus and the electron of the same atom (NA- eA, NB - eB)
2. Between the nucleus and the electron of different atoms (NA- eB, NB - eA)
Repulsive forces −
1. Between the electrons of the two atoms (eA - eB)
2. Between the nuclei of the two atoms (NA - NB)
Experimentally, the attractive forces are found to be stronger than the repulsive ones. Therefore, the two atoms approach each other, and the potential energy also drops during the process.
When the two forces balance each other, a state of minimum energy is attained. At this point, the two atoms are said to be bonded. The distance between them is called the bond length (in case of H-atoms, the bond length is equal to 74 pm).
Also, during bond formation, energy is released. This released energy is called bond enthalpy. The minimum of the given curve represents bond energy, which is 435.8 kJ for hydrogen.
It can be observed in the above figure that when the two H-atoms form a bond, they overlap each other slightly. This represents overlapping of atomic orbitals. (Hence, this theory introduces the concept of atomic orbitals, stating that covalent bonds are formed by the overlapping of singly filled atomic orbitals containing electrons having opposite spins.)
More the extent of the overlapping, stronger will be the bond.
Overlapping of Atomic Orbitals
The overlap between various atomic orbitals can be positive, negative or zero. The given figure shows the vari...