Quantum Mechanical Model - Lesson 1

Quantum Mechanical Model

• Sublevels: Four types - S, P, D, F, each allowing different maximum electron capacities (S: 2, P: 6, D: 10, F: 14).
• Orbitals: Regions of space with high probabilities of finding an electron.
• Orbital Diagrams: Show energy levels, sublevels, and the arrangement of electrons; lowest to highest energy order applies (Bohr-Rutherford principle).
• Principles for Electron Placement:
  • Aufbau Principle: Electrons fill the lowest energy levels first.
  • Pauli Exclusion Principle: A maximum of 2 electrons can occupy one orbital.
  • Hund's Rule: Each orbital at the same energy level gets one electron before any orbital gets a second.

Electron Configurations

• Anions: Negative ions formed by adding electrons.
• Cations: Positive ions created by removing electrons from the highest energy level.
• Condensed Electron Configuration: Uses noble gas symbols to represent inner electrons; only outer electrons shown for reactivity insights.
• Multiple Ion Charges: Larger atoms have complex configurations, allowing multiple electron removal possibilities.

Quantum Numbers

• Bohr Model Limitations: Failed to explain multi-electron atom behaviors; electrons do not follow fixed paths.
• Heisenberg Uncertainty Principle: Position and velocity of an electron cannot be simultaneously measured due to their small size.
• Wave-Particle Duality (De Broglie): Electrons exhibit both wave-like and particle-like properties.
• Schrodinger's Work: Electron motion and energy characterized by complex equations.

Covalent Bonds

• Types of Bonds:
  • Sigma Bonds: Formed by head-on overlap of p or s orbitals; first bond type in a bonding situation.
  • Pi Bonds: Result from the parallel overlap of two p orbitals; less stable than sigma bonds.
• VSEPR Theory: Shapes of molecules determined by the repulsion between valence electrons; lone pairs take more space.

Hybridization Theory

• Describes mixing of atomic orbitals to form new hybrid orbitals during covalent bond formation.
• Aligns with VSEPR in predicting the shapes of molecules based on minimizing electron pair repulsion.

Polarity and Intermolecular Forces

• Electronegativity: The ability of an atom to attract electrons in a bond; differences lead to bond polarity.
• Molecular Polarity: Overall charge distribution in molecules; not all molecules with polar bonds are polar.
• Intermolecular Forces: Forces between molecules, critical in determining state and properties, including melting and boiling points.
• Types:
  • Dispersion Forces: Weak attractions due to temporary dipoles in nonpolar molecules; stronger in larger, more electron-rich molecules.
  • Dipole-Dipole Forces: Stronger interactions between polar molecules due to permanent charges.
  • Hydrogen Bonds: A strong type of dipole-dipole attraction involving H bonded to N, O, or F; critical for properties of water.

Covalent, Ionic and Metallic Bonding

• Covalent Bonds:
  • Nonpolar: Electrons shared equally.
  • Polar: Electrons shared unequally, enhancing overall polarity.
• Metallic Bonds: Characterized by delocalized electrons, leading to conductivity, malleability, and ductility.
• Ionic Bonds: Formed between metals and nonmetals; involve electron transfer, producing ions that create a lattice structure.
• Network Solids: Continuous networks of covalent bonds, as seen in allotropes of carbon (diamond, graphite, fullerene, carbon nanotubes).

Properties of Metals

• Malleability and Ductility: Metals can withstand deformation due to the ability of atomic layers to slide without breaking bonds.
• Electrical and Thermal Conductivity: High due to freely moving delocalized electrons.
• Hardness: Increases with stronger metallic bonds; affected by atomic size and electron density.

Alloys

• Mixtures of metals to improve properties; e.g., steel (iron and carbon) used for its strength.
• Types of Alloys:
  • Substitutional: Atoms replace other atoms.
  • Interstitial: Smaller atoms fit into the gaps of a metal lattice.