Tracovalency

A carbon atom has a total of six electrons occupying the first two shells, i.e., the K-shell has two electrons and the L-shell has four electrons. This distribution indicates that in the outermost shell there are one completely filled 's' orbital and two half-filled 'p' orbitals, showing carbon to be a divalent atom. But in actuality, carbon displays tetravalency in the combined state. Therefore, a carbon atom has four valence electrons. It could gain four electrons to form C^4- anion or lose four electrons to form C⁴⁺ cation. Both these conditions would take carbon far away from achieving stability by the octect rule. To overcome this problem carbon undergoes bonding(π bond)  by sharing its valence electrons. This allows it to be covalently bonded to one, two, three or four carbon atoms or atoms of other elements or groups of atoms. Let us see how carbon forms the single, double and triple bonds in the following examples.

Methane Molecule

Carbon Dioxide Molecule

The electronic configurations of carbon and oxygen are:

Acetylene Molecule

Catenation

The property of self-linking with atoms of the same element is termed Catenation. Carbon has a unique property of linking itself to other carbon atoms to give open chain or/and cyclic structures. Catenation is favored by atoms where atom to atom covalent bond is quite strong.

In carbon, C-C bond energy is very high (347.3 kJ mol-1) causing catenation. Further, the carbon atom due to its tetravalency, can be bonded to two, three or four carbon atoms by forming single and multiple bonds. Therefore, chains of carbon atoms may be linear, branched or cyclic. For example,

Catenation is responsible for the existence of a large number of organic compounds.

Effect of Hybridisation

Hybridisation influences the bond length and bond enthalpy (strength) in organic compounds. The sp hybrid orbital contains more s character and hence it is closer to its nucleus and forms shorter and stronger bonds than the sp³ hybrid orbital. The sp² hybrid orbital is intermediate in s character between sp and sp³ and, hence, the length and enthalpy of the bonds it forms, are also intermediate between them. The change in hybridisation affects the electronegativity of carbon. The greater the s character of the hybrid orbitals, the greater is the electronegativity. Thus, a carbon atom having an sp hybrid orbital with 50%  s character is more electronegative than that possessing sp²  or sp³ hybridised orbitals. This relative electronegativity is reflected in several  physical and chemical properties of the molecules concerned.

Some Characteristic Features of  Bonds

Strength of Bonds with Other Elements

The C-C and C-H bonds in organic compounds are very strong. During chemical reactions, these bonds generally do not break easily. However other atoms or groups such as Cl, OH, etc., attached to the carbon atom may get replaced easily. For example, aqueous KOH reacts with haloalkanes to give the corresponding alcohol.

In a π (pi) bond formation, parallel orientation of the two p orbitals on adjacent atoms is necessary for a proper sideways overlap.Thus, in H2C=CH2    molecule all the atoms must be in the same plane. The p orbitals are mutually parallel and both the p orbitals are perpendicular to the plane of the molecule. Rotation of one CH2 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. The electron charge cloud of the π  bond is located above and below the plane of bonding atoms. This results in the electrons being easily available to the attacking reagents. In general, π  bonds provide the most reactive centres in the molecules containing multiple bonds.