Valence Bond Theory:
The valence bond theory is based on the knowledge of atomic orbitals and electronic configurations of elements overlap criteria of atomic orbitals and stability of molecule. The basic assumptions of valence bond theory are:
i) Atoms do not lose their identity even after the formation of the molecule.
ii) The bond is formed due to the interaction of only the valence electrons as the two atoms come close to each other. The inner electrons do not participate in the bond formation.
iii) During the formation of bond, only the valence electrons from each bonded atom lose their identity. The other electrons remain unaffected.
iv) The stability of bond is accounted by the fact that the formation of bond is accompanied by release of energy. The molecule has minimum energy at a certain distance between the atoms known as internuclear distance. Larger the decrease in energy, stronger will be the bond formed.
To understand the concept more clearly, let us consider the formation of H2 molecule.
Valence Bond Treatment for H2 Molecule:-
Consider two hydrogen atoms A and B approaching each other having nuclei HA and HB and the corresponding electrons eA and eB respectively. When the two atoms are at large distances from each other, no interaction between the two atoms takes place. At this stage, the total energy of the system is the sum of the energies of the individual atoms. When the two atoms come closer, new attractive and repulsive forces begin to operate. Besides the attraction of the nucleus of one atom for its own electrons i.e. HA—eA and HB—eB, the following attractive and repulsive forces start operating:
i) Attractive forces operate between electron of atom A (eA) and nucleus of B (HB) and electron of atom B (eB) and nucleus of A (HA). These two new attractive forces are shown in Fig. (a)
ii) Repulsive forces are present between the nuclei HA—HB and electrons of two atoms eA—eB. These two forces are shown in Fig. (b)
Thus, the forces operating in the molecule are:
Attractive:
i) Nucleus of one atom and its own electron HA — eA and HB — eB.
ii) Nucleus of one atom and electron of other atom: HA — eB and HB — eA.
Repulsive:
i) Electrons of two atoms: eA—eB
ii) Nuclei of two atoms: HA—HB.
The attractive forces tend to bring the atoms closer while repulsive forces tend to push them apart. It has been observed experimentally that the magnitude of the new attractive forces is more than the new repulsive forces. As a result, the two atoms approach each other (Fig. 4.19) and the potential energy of the system decreases. This situation has been shown in Fig. 4.19 (b) where the two atoms are close together. As the two atoms come closer and closer, the system becomes more and more stable due to decrease of energy. Ultimately, a stage is reached where the total forces of attraction balance the forces of repulsion and the system acquires minimum energy shown in Fig. 4.19 (c). At this stage, the two hydrogen atoms are said to be bonded together to form a stable molecule and the distance (r0) between the atoms is known as bond length. For hydrogen molecule, the distance between two hydrogen atoms corresponding to minimum energy is 74 pm. The potential energy of the system when two atoms are brought closer and closer is represented in Fig 4.20. This is called potential energy diagram. In this diagram, when the two atoms are for apart (stage a), there is no attractive or repulsive interactions between them and the potential energy of the system (isolated atoms) is assumed to be zero.
Thus, when the bond is formed, energy is released and therefore the hydrogen molecule is more stable than the individual hydrogen atoms. That is,
H + H ——– H + 435.8 kJ mol-1
This energy corresponding to minimum in the curve is called bond energy.
Conversely, when one mole of H2 molecules is dissociated to hydrogen atoms, 435.8 kJ of energy is needed.
Type of Overlapping and Nature of Covalent Bonds:
Depending upon the type of overlapping, the covalent bonds may be divided into two types
- Sigma (σ) bond
- Pi (π) bond
Sigma (σ) bond: – This type of covalent bond is formed by the end to end overlap of bonding orbitals along the internuclear axis. The overlap is known as head on overlap or axial overlap. The sigma bond is formed by any one of the following types of combinations of atomic orbitals:
1) s-s overlapping: – In this type, two half-filled s-orbitals overlap along the internuclear axis as shown below.
ii) s-p overlapping: – This type of overlapping occurs between the half-filled s-orbital of one atom and half-filled p-orbital of the other atom.
iii) p-p overlapping: – This type of overlapping occurs between half-filled p-orhitals of the two approaching atoms.
Pi (π) bond: – This type of covalent bond is formed by the sidewise overlap of the half-filled atomic orbitals of bonding atoms. Such an overlap is known as sidewise or lateral overlap. The atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbital obtained as a result of sidewise overlap consists of two saucer type charged clouds above and below the plane of the participating atoms.
Strength of sigma and pi-bonds: – The strength of a covalent bond depends upon the extent of overlapping of atomic orbitals forming the bond. During the formation of a sigma bond, the overlapping of orbitals takes place to a larger extent. On the other hand, during the formation of a pi-bond, the overlapping occurs to a smaller extent. Therefore, a sigma bond is stronger than a pi-bond.
It is interesting to note that pi-bond between two atoms is formed in addition to a sigma bond. It is always present in the molecules having multiple bonds (double or triple covalent bond).
Note:
1) Bond Energy: – Bond energy may be defined as the amount of energy required to break a chemical bond so as to separate the bonded atoms in the gaseous state. Bond energy depends upon the following factors viz. size of the participating atoms, multiplicity of bonds and number of lone pairs of electrons.
2) Bond Length: – Bond length may be defined as the average distance between the centres of nuclei of the two bonded atoms in a molecule. Bond length depends upon the size of the atoms and nature of bonds.
3) Bond Angle: – The bond angle may be defined as the average angle between the lines representing the orbitals containing the bonding electrons.