Why do all four bonds in methane look exactly the same? This simple question puzzled chemists for years and led to the development of one of the most important concepts in chemical bonding—hybridization. While valence bond theory successfully explained bonding in many molecules, it fell short when applied to compounds such as methane (CH₄), in which carbon forms four identical covalent bonds.
The concept of hybridization arose because valence bond theory could not explain bonding in many molecules. The
best example is compounds of carbon like CH4. In CH4,
carbon forms four covalent bonds. According to valence bond theory, each
covalent bond is formed by the overlap of two singly filled atomic orbitals.
The electronic configuration of carbon does not allow it to form four bonds, as shown below.
There are only two singly filled orbitals for carbon. The covalency of
carbon is four, and thus it was suggested that one of the electrons from the 2s
orbital jumps or gets excited to the 2pz orbital, thus creating four
singly filled orbitals, i.e.
Now, if these orbitals combine as such, the four bonds formed in CH4
would be of two different energies because the combining capacity of 2s and 2p
orbitals is different. But experiments showed that all four bonds in CH4 are of the same strength and thus have the same energy. This means that
the four singly occupied orbitals have the same energy. This is possible only when
they mix together, redistributing the total energy and taking up the average
energy. Thus, “The mixing of the atomic orbitals, which are of slightly different energy, of an atom so as to redistribute the energy in such a way that formation of new orbitals of equal energy takes place is called hybridization,” and the new orbitals thus formed are called hybrid orbitals.
The type of hybridization deduces the shape of a particular molecule.
Type of Hybridisation
sp Hybridisation:
In this hybridization, one s orbital intermixes with one p orbital, forming two sp hybrid orbitals. sp hybrid orbitals are arranged at 180 degrees and form a linear shape.
sp2 Hybridisation:
one orbital of s subshell and two orbitals of p subshell with comparable energy are mixed to each other and form three hybrid orbitals of equal energy and shape. These hybrid orbitals are named sp2 hybrid orbitals. The hybrid orbitals form a trigonal shape with an angle of 120 degrees; therefore, the molecule's resultant shape is trigonal.
sp3 hybridisation
when one orbital of s subshell is intermix with three orbitals of p subshell then four hybrid orbitals are formed with same shape and energy. Each orbital is known as an sp3 hybrid orbital and is joined to the others at 109.5 degrees, forming a tetrahedral geometry. The molecule containing a sp3 hybrid central atom is tetrahedral in shape.
sp3d hybridisation
This hybridization results from the intermixing of five orbitals (one "s" orbital, three "p" orbitals, and one "d" orbital) from subshells of comparable energy. In this hybridization, five hybrid orbitals are formed with similar energy and the same shape. These orbitals form a trigonal bipyramid. The bond angles are found to be 120 degrees and 90 degrees.
sp3d2 hybridisation
sp3d2 hybridisation is also known as octahedral hybridisation. This hybridisation results from the intermixing of six orbitals from the s-, p-, and d-subshells. The bond angle between hybrid orbitals is 90 degrees.
sp3d3 hybridisation
The molecule of iodine heptafluoride is pentagonal bipyramidal because the Iodine atom is sp3d3 hybridised in this molecule.
A summary of the hybridization type shape relationship
is given below:
Four simple steps for the determination of Hybridisation
Step 1: Draw the Lewis dot structure (Electron dot structure)
Example
Step 2: Count the number of lone pair and bond pairs at the central atom
Step 3: Find the hybridization from the following table as per the number of lone pairs and bond pairs.
Step 4: Conclusion
The carbon atom in the above molecule is sp3-hybridized, and the bond angle is 109.5 degrees.
