When two hydrogen atoms approach each other, their 1s atomic orbitals interact to form a bonding and an antibonding molecular orbital (MO). A stable covalent bond is formed when the bonding MO is filled with a pair of electrons and the antibonding MO is empty.
Sigma (σ) bonds are strong bonds with a circular cross-section formed by the head-on overlap of two atomic orbitals.
The electronic configuration of atomic carbon implies that carbon should form two bonds. However, it is known that carbon forms four bonds. When carbon is part of an organic structure, it can ‘mix’ the 2s and 2p orbitals of the valence shell in a process known as hybridization. There are three possible types of hybridization – sp3, sp2 and sp hybridization.
A covalent bond binds two atoms together in a molecular structure and is formed when atomic orbitals overlap to produce a molecular orbital – so called because the orbital belongs to the molecule as a whole rather than to one specific atom. A simple example is the formation of a hydrogen molecule (H2) from two hydrogen atoms. Each hydrogen atom has a half-filled 1s atomic orbital and when the atoms approach each other, the atomic orbitals interact to produce two MOs (the number of resulting MOs must equal the number of original atomic orbitals, Fig. 1).
The MOs are of different energies. One is more stable than the original atomic orbitals and is called the bonding MO. The other is less stable and is called the antibonding MO. The bonding MO is shaped like a rugby ball and results from the combination of the 1s atomic orbitals. Since this is the more stable MO, the valence electrons (one from each hydrogen) enter this orbital and pair up. The antibonding MO is of higher energy and consists of two deformed spheres. This remains empty. Since the electrons end up in a bonding MO which is more stable than the original atomic orbitals, energy is released and bond formation is favored. In the subsequent discussions, we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape, but it is important to realize that antibonding molecular orbitals also exist.
The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond: σ bonds have a circular cross-section and are formed by the head-on overlap of two atomic orbitals. This is a strong interaction and so sigma bonds are strong bonds. In future discussions, we shall see other examples of σ bonds formed by the interaction of atomic orbitals other than the 1s orbital.
Atoms can form bonds with each other by sharing unpaired electrons such that each bond contains two electrons. We identified that a carbon atom has two unpaired electrons and so we would expect carbon to form two bonds. However, carbon forms four bonds! How does a carbon atom form four bonds with only two unpaired electrons?
So far, we have described the electronic configuration of an isolated carbon atom. However, when a carbon atom forms bonds and is part of a molecular struc-ture, it can ‘mix’ the s and p orbitals of its second shell (the valence shell). This is known as hybridization and it allows carbon to form the four bonds which we observe in reality.
There are three ways in which this mixing process can take place.
· the 2s orbital is mixed with all three 2p orbitals. This is known as sp3 hybridization;
· the 2s orbital is mixed with two of the 2p orbitals. This is known as sp2 hybridization;
· the 2s orbital is mixed with one of the 2p orbitals. This is known as sp hybridization.
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