Chap_03B-Modern_Atomic_Theory_updated_10_3_2014

Modern Atomic Theory

Chapter Overview

• Chapter covers the evolution of atomic theory, focusing on electron arrangements, Bohr model, quantum mechanical model, electron configuration and periodic properties.

Chapter Outline

• Bohr Model of Atom

• Quantum Mechanical Model of Atom

• Electron Configuration

• Electron Configuration & Periodic Table

• Abbreviated Electron Configuration

• Periodic Properties

Electron Arrangement in Atoms

• Nucleus and Electrons: Protons and neutrons are in the dense nucleus; electrons exist in the surrounding space.

• Physical and Chemical Properties: The arrangement of electrons determines the properties of elements.

• Energy Levels: Electrons occupy specific energy levels essential for understanding their behavior.

Bohr Model of Atom

Initial Model by Niels Bohr

• Developed in early 20th century.

• Electrons orbit nucleus at defined distances called energy levels.

Key Characteristics

• Electrons occupy specific energy levels.

• Can jump to higher energy levels by absorbing energy.

• Ground State: Lowest energy level.

• Excited States: Higher energy levels.

• Energy is emitted as light when excited electrons return to ground state.

Limitations

• Cannot fully explain electron behavior in larger atoms.

Quantum Mechanical Model of Atom

Development by Erwin Schrödinger

• Introduced in 1926.

• Predicts electron locations within a probability region called orbitals.

Principal Energy Levels

• Designated by the letter 'n'.

• Electrons can exist at various energy levels (higher n = higher energy).

Sublevels

• Sublevels are designated as s, p, d, and f.

• Each sublevel contains orbitals:

  • s: 1 orbital (2 electrons)

  • p: 3 orbitals (6 electrons)

  • d: 5 orbitals (10 electrons)

  • f: 7 orbitals (14 electrons)

Electron Configuration

Importance of Electron Arrangement

• Similarities in periodic table behaviors stem from electron configurations.

• Electrons occupy orbitals from lowest to highest energy levels: s < p < d < f.

Filling Orbitals

• Each orbital must be singly occupied before double (Hund’s Rule).

Notation Examples

• Complete Notation: Example for Phosphorus (Z=15): 1s² 2s² 2p⁶ 3s² 3p³

• Orbital Notation: Shows electron spins.

Abbreviated Electron Configuration

Streamlining Configurations

• Utilizes noble gas notation to simplify configurations for larger atoms.

• Example: Potassium (K, Z=19): [Ar] 4s¹.

Periodic Properties

Trends in Atomic Properties

• Properties influenced by electron configurations include atomic size, ionization energy, and metallic character.

Atomic Size

• Trends: Increases down a group due to added energy levels; decreases across a period due to increased nuclear charge.

Ionization Energy

• Definition: Energy needed to remove a valence electron.

• Decreases down a group; increases across a period due to stronger nuclear attraction.

Metallic Character

• Ability to lose electrons easily. Decreases across a period and increases down a group.

Examples

Atomic Radius Comparisons

1. Li vs K: K has a larger radius due to more energy levels.

2. K vs Br: K has a larger radius due to less nuclear charge.

3. P vs Cl: P has a larger radius due to less nuclear charge.

Ionization Energy Comparisons

1. Na vs K: Na has a higher ionization energy due to fewer energy levels.

2. Mg vs Cl: Cl has a higher ionization energy due to greater nuclear charge.

3. F vs C: F has the highest ionization energy due to most nuclear charge.

Conclusion

• Understanding these atomic theories and properties is essential for analyzing elemental behaviors and conducting further scientific inquiry.