Hybridization

Covalent Bond: A chemical bond that involves the sharing of electron pairs between atoms.
Understanding through Valence Bond Theory (VBT):
- Focuses on outermost electrons (valence electrons) to describe how atoms bond.
- Considers the concept of overlapping atomic orbitals between different atoms.

Atomic Orbitals
- Describes regions around an atom where electrons are likely to be found.
- Important orbitals include the s and p orbitals.
Half-Filled Orbitals: At least one unpaired electron is present in an orbital, necessary for bond formation.

Example of H₂ (Hydrogen Molecule)
- Two hydrogen atoms come together, each with one electron in a 1s orbital.
- Bond Formation:
  - The half-filled 1s orbitals from each hydrogen atom overlap to form a covalent bond.
  - Resulting molecular orbital is a larger orbital containing the shared pair of electrons.
- Probability Distribution: Electrons are not confined but spread out in space around both nuclei.
Example of H₂S (Hydrogen Sulfide)
- Hydrogen contributions: Two half-filled 1s orbitals from hydrogen.
- Sulfur contributions: Two valence electrons that exist in three p orbitals.
- Appeal of orbital overlap in bonding: 1s from H overlaps with 3p on S.
- This results in two bonds.
- Bond Angle Prediction:
  - Expected Bond Angle = 90 degrees due to perpendicular nature of p orbitals.
  - Experimental Bond Angle = 92 degrees, showing accuracy of predictive model.

Example of Carbon Bonding with Hydrogen:
- Carbon's ground state electron configuration signifies two half-filled orbitals.
- Predicted only two bonds with hydrogen, which does not reflect reality (actual has four bonds).
- Geometry Issues: 90 degrees predicted vs real-world shapes.
Need for Hybrid Orbitals: Required to describe molecules, not just singular atoms.

Creating Hybrid Orbitals:
- Definition: Hybrid orbitals result from the mathematical combination of atomic orbitals (s and p) to form new orbitals optimizing bond formation.
- SP Hybrid Orbitals:
- Formed by mixing one s orbital and one p orbital.
- Two resulting SP orbitals have larger lobes allowing for linear bonding.
Example of Hybridization in BeCl₂ (Beryllium Chloride):
1. Lewis Structure Construction: Linear arrangement predicted by VSEPR theory.
2. Beryllium has full 2s and empty 2p orbital - hybridizing to create two SP orbitals, allowing equivalent bond formation.
- Result: Linear geometry with bond angles of 180 degrees.
Sigma Bonds:
- Formed from end-on overlap of atomic orbitals.
- Specifically between s and p orbitals generates sigma bonds.

Linear Geometry: SP Hybridization
Trigonal Planar Geometry: SP² Hybridization
- Three orbitals will be used to accommodate three surrounding atoms in a plane.
- Hybrid Orbitals Orientation: Points towards corners of a triangle to minimize repulsion.
Tetrahedral Geometry: SP³ Hybridization
- Four hybrid orbitals resulting in a tetrahedron shape.
Trigonal Bipyramidal Geometry: SP³D Hybridization
- Comprised of hybridized s, p, and d orbitals.
Octahedral Geometry: SP³D² Hybridization
- Six orbitals formed for structuring octahedral shapes.

Determines spatial arrangement and types of bonds (sigma and pi).
Pi Bonds: Formed by sideways overlap of unhybridized p orbitals, adding an additional layer to bonding.