USA: +1-585-535-1023

UK: +44-208-133-5697

AUS: +61-280-07-5697

Types of Overlapping and Nature of Covalent Bonds

The formation of a covalent bond involves the overlapping of half-filled atomic orbitals. The covalent bonds can be classified into two different categories depending upon the type of overlapping. These are:

(a) Sigma covalent bond

(b) Pi covalent bond.

(a) Sigma ( σ) bond. This type of covalent bond is formed by the axial overlapping of half-filled atomic orbitals. The atomic orbitals overlap along the inter-nuclear axis and involve end to end or head on overlap. The electron cloud formed as a result of axial overlap is cylindrically symmetrical about inter-nuclear axis. The electrons constituting sigma bond are called sigma electrons. There can be three types of axial overlap among s and p-orbitals as discussed below:

(z) s-s overlap. It involves mutual overlap of half-filled s-orbitals of the atoms approaching to form a bond. The bond formed is called s-s σ bond •

(ii) s-p overlap. It involves mutual overlap of half-filled s-orbital of the one atom with half-filled p-orbital of the other. The bond so formed is called s-p σ bond.

(iii) p-p overlap. It involves mutual overlap of half-filled p-orbitals of the two atoms. The bond so formed is called p-p σ bond.

The s-s, s-p and p-p overlaps have been shown diagrammatically in Fig. 36.1.

(b) Pi (1t) Bond. This type of covalent bond is formed by the lateral or sidewise overlap of the atomic orbitals. The orbital overlap takes place in such a way that their axes are parallel to each other but perpendicular to the internuclear axis. The pi bond consists of two charge clouds above and below the plane of the atoms involved in the bond formation. The electrons involved in the 1t-bond formation are called n-electrons.

SOME CHARACTERISTIC FEATURES OF π – BONDS

A pi (π) bond is constituted by side ways overlap of orbital perpendicular to the internuclear axis, some characteristic features are:

(i) All the atoms directly attached to the carbon atoms of double bond lie in the same plane. For example, in CH2 = CH2 all the six atoms (2 carbon atoms and 4 hydrogen atoms) lie in the same plane.

(ii) Only the unhybridised p-orbitals perpendicular to the plane of the molecule from pi bonds.

(iii) Rotation of one C~ fragment with respect to other interferes with maximum overlap of p-orbitals and, therefore, such rotation about carbon-carbon double bond (C = C) is restricted.

(iv) The electron charge cloud of the 1t-bond is placed above and below the plane of bonding atoms. This results in the electrons being easily available to the attacking reagents.

In general, n-bonds provide the most reactive centres in the molecules containing multiple bonds.

It may be noted that:

(i) Sigma bond is stronger than pi bond. It is because of the fact that overlapping of atomic orbitals can take place to a greater extent during the formation of sigma bond whereas overlapping of orbitals occurs to a smaller extent during the formation of pi bond.

(ii) Pi bond between the two atoms is formed only in addition to a sigma bond. It is because of the fact that the atoms constituting a single bond prefer to form a strong sigma bond rather than a weak pi bond. Thus, pi bond is always present in molecules having multiple bonds, i.e., double or triple bond. In other words, a single bond cannot be a pi bond.

(iii) The shape of molecule is controlled by the sigma framework (orientations of sigma bonds) around the central atom. Pi bonds are superimposed on sigma bonds hence they simply modify the dimensions of the molecule.