Atomic Orbitals and Hybridization

Three-Dimensional Nature of Atomic Orbitals

Atomic orbitals describe regions in space where electrons are likely to be found. These orbitals have distinct three-dimensional shapes that influence chemical bonding and molecular geometry.

• s orbitals: Spherical in shape, centered around the nucleus (e.g., 1s, 2s).

• p orbitals: Dumbbell-shaped, oriented along the x, y, and z axes (e.g., 2px, 2py, 2pz).

• Hybrid orbitals: Formed by the combination of atomic orbitals (e.g., sp3, sp2, sp), and are crucial for understanding bonding in organic molecules.

Example: The carbon atom in methane (CH4) uses sp3 hybrid orbitals to form four equivalent bonds with hydrogen.

Ionic and Covalent Bonding

Ionic Bonds (Brief Overview)

While not the main focus in organic chemistry, ionic bonds are important for comparison. They result from electrostatic attraction between oppositely charged ions.

• Cations (+) and anions (−) attract each other.

• Strength increases with larger charges and smaller ion sizes.

• High melting points (e.g., NaCl ~ 800°C).

• Soluble in water, but generally insoluble in organic solvents.

Example: Sodium chloride (NaCl) is a classic ionic compound.

Covalent Bonding in Dihydrogen (H2)

Valence Bond Model

The valence bond model explains covalent bonding as the overlap of atomic orbitals, resulting in a shared pair of electrons.

• Each hydrogen atom contributes one 1s electron.

• In-phase overlap of two 1s orbitals forms a bonding orbital containing two electrons.

Electrostatic View: The electron density is concentrated between the nuclei, holding the atoms together.

Molecular Orbital Model

The molecular orbital (MO) model describes bonding as the combination of atomic orbitals to form molecular orbitals that extend over the entire molecule.

• Bonding MO: Constructive overlap (in-phase) lowers energy and stabilizes the molecule.

• Antibonding MO: Destructive overlap (out-of-phase) creates a node and raises energy.

Energy Diagram:

• Bonding orbital is lower in energy than the original atomic orbitals.

• Antibonding orbital is higher in energy.

Molecular Geometry and VSEPR Theory

Electron Pair Repulsion and Molecular Shape

The shape of a molecule is determined by the repulsion between electron pairs (bonded and lone pairs) around a central atom, as described by the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• Bonded pair–bonded pair repulsion: Least repulsive.

• Lone pair–bonded pair repulsion: More repulsive.

• Lone pair–lone pair repulsion: Most repulsive.

Example: In methane (CH4), the four bonding pairs arrange themselves tetrahedrally (109.5°) to minimize repulsion.

Common Molecular Geometries

Different molecules adopt characteristic shapes based on the number of electron pairs around the central atom.

• Tetrahedral (e.g., CH4): 109.5° bond angles.

• Trigonal planar (e.g., formaldehyde, H2CO): 120° bond angles.

• Linear (e.g., carbon dioxide, CO2): 180° bond angles.

Example: Formaldehyde (H2CO) is trigonal planar due to three regions of electron density around the central carbon.

Table: Summary of Common Molecular Geometries

AXn notation: A = central atom, X = surrounding atoms, n = number of X atoms

Additional info: The slides reference the 10th edition of "Organic Chemistry" by T. W. Graham Solomons, which is a standard textbook for college-level organic chemistry. The content aligns with foundational topics in atomic structure, bonding, and molecular geometry, which are essential for understanding organic molecules and their reactivity.