ch03

Chapter 3: The Quantum Model of the Atom and Periodicity

3.1 A Brief Exploration of Light

• Dual Nature of Light: Light exhibits both wave-like and particle-like properties.

• Energy Relationships:

  • Energy of light (E) is related to frequency (ν) and wavelength (λ).

  • Important calculations relating energy, frequency, and wavelength are possible.

3.2 Bohr Theory of the Atom

• Bohr Model: Describes energy levels in atoms, explaining light emission and absorption by gaseous atoms.

• Energy Transitions: Can calculate energy or wavelength for specific transitions in hydrogen atoms.

3.3 Wave-Particle Duality of Matter

• De Broglie Hypothesis: Matter also has wave-like properties; wavelength relates to mass and velocity.

• Heisenberg's Uncertainty Principle: States that position and momentum (velocity) of electrons cannot be simultaneously known with arbitrary precision.

3.4 Orbitals and Quantum Numbers

• Quantum Numbers: Identify the energy state and position of electrons in atoms.

  • Principal Quantum Number (n): Indicates energy level (shell).

  • Angular Momentum Quantum Number (ℓ): Indicates subshell shape.

  • Magnetic Quantum Number (mℓ): Indicates orbital orientation.

  • Spin Quantum Number (ms): Indicates electron orientation within an orbital.

3.5 Shape of Orbitals

• Orbital Types:

  • s Orbitals: Spherical shape; can hold 2 electrons.

  • p Orbitals: Dumbbell shape; contains three orbitals (total of 6 electrons).

  • d and f Orbitals: More complex shapes with more orbitals, increasing electron capacity.

3.6 Orbital Diagrams

• Energy-Level Diagrams: Visual representation showing the arrangement of orbitals and electron distribution.

• Aufbau Principle: Refers to filling orbitals from lowest to highest energy.

• Hund’s Rule: Electrons occupy degenerate orbitals singly before pairing.

3.7 Electron Configurations

• Electron Configuration Notation: Represents electron arrangement in atoms by specifying subshells and their electron counts.

• Valence Electrons: Outermost electrons involved in chemical bonding; determine an element's reactivity and placement in groups.

3.8 Valence Electrons

• Valence vs Core Electrons:

  • Valence: In outermost shell.

  • Core: In inner shells.

• Valence electrons dictate periodicity in chemical properties.

3.9 Atomic and Ionic Sizes

• Atomic Radius Trends:

  • Decreases across a period, increases down a group.

  • Influenced by effective nuclear charge (Zeff): the net positive charge felt by the valence electrons.

• Ionic Size Trends:

  • Cations: Smaller than their neutral atoms due to loss of electron-electron repulsion.

  • Anions: Larger than their neutral atoms due to increased electron-electron repulsion.

3.10 Ionization Energy and Electron Affinity

• Ionization Energy (IE): Energy required to remove an electron from a gaseous atom. Generally increases across a period and decreases down a group.

• Electron Affinity (EA): Energy change when an electron is added to a gaseous atom; often exothermic for non-metals, becoming more negative across periods.