LM_Fall2024_CHEM1151_Topic_11 - Metallic Bonding and Band Theory

Periodic Structures: Metals form structures similar to ionic crystals with identical atoms.
Key Properties:
- Plasticity: Ability to be deformed without breaking.
- Electrical Conductivity: Metals conduct electricity well due to free-moving electrons.
- Thermal Conductivity: Efficient heat transfer through mobile electrons.
- Luster: Reflective qualities caused by delocalized electrons that absorb and re-emit light.
Metallic Bonding: Atoms in metals are held together by metallic bonding, where electrons are shared among all atoms.

Electron Behavior: Metals have low electronegativity (EN), allowing valence electrons to be shared rather than localized.
Structure: Fixed positive ions surrounded by a "sea" of delocalized electrons.
Macroscopic Properties: Explains various physical characteristics of metals such as malleability, ductility, and thermal/electrical conductivity.

Mechanism: As heat energy increases, electrons gain kinetic energy, colliding with other electrons, thus transferring energy and raising temperature.

Directionality of Electrons: A chaotic electron motion becomes organized by an electric field, allowing for efficient conduction of electricity.

Comparison with Ionic Compounds: Unlike ionic compounds that are hard and brittle, metals are malleable (can be flattened) and ductile (can be drawn into wires).
Mechanism: The presence of delocalized electrons allows cations to shift position smoothly under applied force without fracturing.

Photon Absorption: Electrons absorb photons and get excited, later re-emitting them as they return to a lower energy state.
Shininess: This interaction with visible light gives metals their characteristic shiny appearance.

Basis of Theory: A more comprehensive explanation of molecular properties compared to valence bond theory.
- Formation: Molecular orbitals (MOs) result from the combination of atomic orbitals (AOs).
- Orbitals per Atom: The number of MOs equals the number of AOs combined.
Phase of Atomic Orbitals: The sign of the wavefunction indicates atomic orbital phases, essential for bonding.
- Bonding Orbitals: Formed from orbitals with the same phase (constructive interference).
- Antibonding Orbitals: Formed from orbitals with opposite phases (destructive interference), designated with an asterisk (*).

σ and π Orbitals: Define electron density distribution.
- σ Orbitals: Cylindrical electron density between nuclei, with no nodal planes.
- π Orbitals: Electron density above and below a nodal plane containing the nuclei.
Molecular Orbital Diagrams: Represent the combination of AOs to form MOs, illustrating electrons' arrangement.

Bond Order Calculation: Defined as Bond Order = (Number of Bonding Electrons - Number of Antibonding Electrons) / 2.
- Stability Indicator: A bond order greater than zero indicates stability.
Example of H2 Molecule: Bond order of 1 signifies a stable single bond.

Band Theory:
- Band Gap: The energy separation between the conduction band and valence band.
  - Insulators have large band gaps (typically >3.5 eV).
- Conductors: Metals have overlapping valence and conduction bands, allowing free electron movement.

Band Gap: Semiconductors possess a small band gap, allowing some electrons to be thermally excited into the conduction band.
Types of Semiconductors:
- Elemental Semiconductors (e.g., Si, Ge): Contain single type atoms.
- Compound Semiconductors (e.g., GaAs): Formed from two or more types of atoms.

Intrinsic vs. Extrinsic:
- Intrinsic: Pure form of semiconductor without impurities.
- Extrinsic: Doped semiconductors that have modified conductivity.
Doping Process:
- N-type Semiconductors: Introduced with extra valence electrons (e.g., Si doped with P).
- P-type Semiconductors: Created by introducing atoms with fewer valence electrons (e.g., Si doped with Ga).
- Significance: A single dopant can vastly improve conductivity (up to 100,000 times).