Hybridization note

Definition: Hybridization is the process of mixing different orbitals of the same atom to form a new set of equivalent orbitals known as hybrid orbitals.
Purpose: Used to ascertain information on bonding and shape of the molecule, such as linear, triangular, or tetrahedral shapes.
Characteristics:
- The mixing of atomic orbitals allows the formation of directional hybrid orbitals.
- High electron density lobes point in specific directions at correct angles, leading to more effective overlap compared to s and p orbitals, resulting in stronger bonds.

Atoms involved: One carbon atom (C) and four hydrogen atoms (H)
Electronic Configuration:
- C: 1s² 2s² 2p
- H: 1s¹
Process:
1. The 2s orbital and three 2p orbitals of carbon mix to form four equivalent sp³ hybrid orbitals.
2. These sp³ orbitals overlap with 1s orbitals of four hydrogen atoms, forming stronger bonds due to high electron density overlap.
Characteristics of sp³ Hybrid Orbital:
- Equivalent to the other three with 25% s-character and 75% p-character.
- Tetrahedral structure with bond angle of 109° 28'.

Process:
1. Two p orbitals and one s orbital mix to give three coplanar sp² hybrid orbitals directed at 120° (planar triangle).
2. Boron must be excited to have three unpaired electrons to form three covalent bonds.
Electronic States:
- Ground state: 1s² 2s² 2p
- Excited state: 1s² 2s² 2p¹ 2p¹ 2p¹
Bond Formation:
- Empty p orbital can accept an electron from a chlorine atom to form a dative bond.
- Three equivalent sp² hybrid orbitals combine with three atomic orbitals of chlorine atoms to form six molecular orbitals (3 bonding and 3 antibonding).

Electronic Configuration:
- Be: 1s² 2s²
- Excited state: 1s² 2s² 2p
Geometric Configuration:
- Two equivalent sp hybrid orbitals directed at 180°, resulting in a linear structure.
Bond Formation:
- Combine with the atomic orbitals of chlorine to create four molecular orbitals (2 bonding and 2 antibonding).

Definition: Formation of hybrids with more s-character or more p-character to adjust bond angles.
Ammonia (NH₃): Requires more p-character to achieve bond angles of 107°.
Water (H₂O): Achieved bond angle of 105° due to significant lone pair repulsion.
Expected Angles without mixing: 90° for the p orbital bonding.

General Principle: More than two atoms lead to molecular shape consideration. The shape arises from minimizing total energy in the ground state.
Sidgwick and Powell Theory (1940):
  1. Shapes can be approximated based on electron pairs around the central atom.
  2. Bond pairs and lone pairs occupy space and repel each other.
  3. Orientation minimizes electron pair repulsion.

Gillespie and Nyholm Modifications (1957):
  1. Molecular shape determined by repulsions between all electron pairs in the valence shell.
  2. Lone pairs cause distortion in bond angles. The order of repulsion strength is:
     - Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
  3. Magnitude of repulsions varies with electronegativity differences.
  4. Double bonds have more repulsion than single bonds; triple bonds more than double bonds.

Electron Pairs: Hybridization and corresponding geometries and angles.
  - 2 pairs:
    - Geometry: Linear
    - Hybridization: sp
    - Bond angle: 180°
  - 3 pairs:
    - Geometry: Trigonal Planar
    - Hybridization: sp²
    - Bond angle: 120°
  - 4 pairs:
    - Geometry: Tetrahedral
    - Hybridization: sp³
    - Bond angle: 109° 28'
  - 5 pairs:
    - Geometry: Trigonal Bipyramidal
    - Hybridization: sp³d
    - Bond angles: 120° and 90°
  - 6 pairs:
    - Geometry: Octahedral
    - Hybridization: sp³d²
    - Bond angle: 90°
  - 7 pairs:
    - Geometry: Pentagonal Bipyramidal
    - Hybridization: sp³d³
    - Bond angles: 72° and 90°

Examples:
- BeCl₂, H-C≡C-H, O=C=O
- Multiple bonds are treated as equivalent to a single electron pair (single covalent bond).

Isoelectronic Molecules:
  - CH₄: Regular tetrahedron, bond angle 109° 28'.
  - NH₃: Three bond pairs and one lone pair lead to a bond angle of 107° 48'.
  - H₂O: Two lone pairs and two bond pairs lead to a bond angle smaller than ideal tetrahedral (104° 27').

Examples:
- NH₃ vs NF₃ or H₂O vs F₂O
- Higher electronegativity in fluorine leads to decreased bond angles due to higher pull on bonding electrons, resulting in broader spread of bond pairs.
Bond Angles:
- H-N-H: 107°
- F-N-F: 102°
General Observations:
- Decrease in bond angles corresponds to lower electronegativity or a more central bonding characteristics, affecting repulsion among bonds.