sp2 HYBRIDIZATION

Orbital hybridization

In chemistry, hybridization is the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties. Although sometimes taught together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and hybridization are in fact not related to the VSEPR model.

History

ChemistLinus Pauling first developed the hybridization theory in 1931 in order to explain the structure of simple molecules such as methane (CH₄) using atomic orbitals. Pauling pointed out that a carbon atom forms four bonds by using one s and three p orbitals, so that "it might be inferred" that a carbon atom would form three bonds at right angles (using p orbitals) and a fourth weaker bond using the s orbital in some arbitrary direction. In reality however, methane has four bonds of equivalent strength separated by the tetrahedral bond angle of 109.5°. Pauling explained this by supposing that in the presence of four hydrogen atoms, the s and p orbitals form four equivalent combinations or hybrid orbitals, each denoted by sp³ to indicate its composition, which are directed along the four C-H bonds. This concept was developed for such simple chemical systems, but the approach was later applied more widely, and today it is considered an effective heuristic for rationalizing the structures of organic compounds. It gives a simple orbital picture equivalent to Lewis structures. Hybridization theory finds its use mainly in organic chemistry.

sp² HYBRIDIZATION

Boron trichloride (BCl₃)

* The electronic configuration of 'B' in ground state is 1s² 2s² 2p¹ with only one unpaired electron. Since the formation of three bonds with chlorine atoms require three unpaired electrons, there is promotion of one of 2s electron into the 2p sublevel by absorbing energy.Thus Boron atom gets electronic configuration: 1s² 2s² 2p_x¹2p_y¹. 

However to account for the trigonal planar shape of this BCl₃ molecule, sp² hybridization before bond formation was put forwarded.* In the excited state, Boron undergoes sp² hybridization by using a 2s and two 2p orbitals to give three half filled sp² hybrid orbitals which are oriented in trigonal planar symmetry. * Boron forms three σ_sp-p bonds with three chlorine atoms by using its half filled sp² hybrid orbitals. Each chlorine atom uses it'shalf filled p-orbital for the σ-bond formation.

* Thus the shape of BCl₃ is trigonal planar with bond angles equal to 120^o.

Ethylene (C₂H₄)

* During the formation of ethylene molecule, each carbon atom undergoes sp² hybridization in its excited state by mixing 2s and two 2p orbitals to give three half filled sp² hybrid orbitals oriented in trigonal planar symmetry. 

There is also one half filled unhybridized 2p_z orbital on each carbon perpedicular to the plane of sp² hybrid orbitals. 

* The carbon atoms form a σ_sp2_-sp2 bond with each other by using sp² hybrid orbitals. 

A π_p-p bond is also formed between them due to lateral overlapping of unhybridized 2p_z orbitals. 

Thus there is a double bond (σ_sp2_-sp2 &  π_p-p) between two carbon atoms. 

* Each carbon atom also forms two σ_sp2_-s bonds with two hydrogen atoms. 

* Thus ethylene molecule is planar with ∠HCH &∠HCC bond angles equal to 120^o. 

* All the atoms are present in one plane.