Carbon atoms cannot form four bonds in their unhybridized state. Carbon must hybridize to form four covalent bonds. The hybridization states of carbon are sp3, sp2, and sp.
When carbon atoms are bonded to other atoms, they use hybrid orbitals. A hybrid orbital is an electron orbital that forms when two atomic orbitals combine to form a covalent bond. In its atomic state, carbon has the valence electron configuration 1s22s22px12py1, with only two half-filled orbitals.
To form four single bonds, carbon mixes the s and p orbitals into four sp3 orbitals, all with equal energy, exhibiting 25% s character and 75% p character. The sp3 orbitals have one large lobe and one small lobe, making the electron density greater on the larger-lobed side of the orbital.
Hybridization of Carbon Orbitals
Methane, CH4, contains a carbon atom with four hybridized orbitals. Each orbital contains an unpaired electron singly bonded to each hydrogen atom. The molecule forms a tetrahedral shape.
Tetrahedral Shape of Methane
When carbon forms a double bond (bonds to three other atoms), the carbon will form three sigma (σ) bonds using sp2 hybridization, which combines one s orbital with two p orbitals to form three new sp2 orbitals. One p orbital is left unhybridized. The two carbon atoms of ethylene bond to each other and to two hydrogen atoms using those three sp2 orbitals. The remaining two p orbitals overlap between the carbon atoms to form a double bond, which is shorter and stronger than a carbon-carbon single bond.
When carbon forms a triple bond or two double bonds (bonds to two other atoms), as in acetylene (C2H2), two hybridized sp orbitals are created, and two unhybridized p orbitals remain. The carbons each form a bond to hydrogen with one sp hybrid orbital. They bond to each other with one sp orbital and two p bonds.
In summary, carbon with all single bonds has sp3 hybridization. Carbon with one double bond has sp2 hybridization. Carbon with two double bonds or one triple bond has sp hybridization.