What are the hybrid orbitals for CH4?
The correct answer and explanation is :
In methane (CH₄), the carbon atom forms four covalent bonds with four hydrogen atoms. To explain the bonding in methane, we need to understand the concept of hybridization.
Hybridization of Carbon in CH₄:
Carbon has an electron configuration of 1s² 2s² 2p². In its ground state, carbon has two unpaired electrons in the 2p orbitals, which is not enough to form four bonds as in methane. To form four equivalent bonds, carbon undergoes hybridization.
- Promotion of Electrons: One of the electrons from the 2s orbital is promoted to an empty 2p orbital. This results in the carbon atom having four unpaired electrons, one in the 2s orbital and three in the 2p orbitals.
- Hybridization: These four orbitals (one 2s and three 2p) then combine to form four new equivalent orbitals. This is called sp³ hybridization. The resulting hybrid orbitals are degenerate, meaning they all have the same energy, and they are directed toward the corners of a tetrahedron, with an angle of 109.5° between them.
- Bond Formation: These sp³ hybrid orbitals each overlap with the 1s orbital of a hydrogen atom, forming a covalent bond. In total, carbon forms four bonds with four hydrogen atoms, creating the methane molecule (CH₄).
Key Points of sp³ Hybridization in CH₄:
- Geometry: The sp³ hybrid orbitals are arranged in a tetrahedral geometry, which minimizes electron repulsion and gives the molecule its symmetrical shape.
- Bonding: Each sp³ orbital on carbon overlaps with a 1s orbital of hydrogen, forming strong sigma (σ) bonds.
Thus, the hybridization of the carbon atom in methane (CH₄) is sp³, resulting in a tetrahedral shape with four equivalent sigma bonds.