Shielding and Effective Nuclear Charge

Introduction

• Shielding and effective nuclear charge are interconnected concepts.
• They help us understand atomic properties influenced by positive (protons) and negative (electrons) particles.
• Focus on valence electron shell where chemistry occurs.

Forces in Atoms

• Attractive Force: Between negative electrons and positive protons.
• Repulsive Force: Electrostatic repulsion among electrons.
• All atoms (except hydrogen) have multiple protons and electrons.
• Examining repulsion and attraction helps analyze atomic behavior and properties.

Second Period Elements

• Lithium (Li): 3 protons, 1 valence electron, 2 core electrons.
  • Protons are represented as blue arrows (attraction).
  • Core electrons have a repulsive effect on valence electrons (violet arrows).
  • Valence electrons feel the attractive force of only one proton after accounting for repulsion.

Definitions

• Shielding: The repelling effect of inner electrons on valence electrons; reduces attraction from the nucleus.

• Effective Nuclear Charge (Z_EFF): The net attraction felt by valence electrons, calculated as:

  Z_EFF = Total Protons (Z) - Shielding Electrons (S)

Beryllium (Be)

• Nuclear charge: 4+ (one more than lithium).
• 2 core electrons, 2 valence electrons.
• Z_EFF for beryllium = 4 - 2 = 2 (greater than lithium).

Fluorine (F)

• Nuclear charge: 9+.
• Z_EFF = 9 - 2 = 7.
• Atomic Size Trend: Atoms decrease in size across the period due to increased nuclear charge pulling electrons inward.

Transition to Period 3

• Compare Neon (Ne) and Sodium (Na):
  • Neon: 10 protons, 2 core electrons, Z_EFF = 8.
  • Sodium: Higher energy shell (n=3), 1 valence electron, greater distance from nucleus.
  • Core electrons increase from 2 to 10, leading to reduced Z_EFF = 1.

Group Trends

• Compare Lithium and Sodium:
  • Both have Z_EFF = 1, but sodium's valence electron is further from the nucleus.
  • Pattern observed: Atomic size increases down a group due to increased shielding and energy levels.

Periodicity in Atomic Size

• Across a Period: Decrease in atomic size due to increased protons and constant shielding.
• Down a Group: Increase in atomic size due to added energy levels and increased shielding.

Ionization Energy (IE)

• Energy required to remove an electron.
• IE increases across a period due to closer valence electrons to the nucleus.
• IE decreases down a group due to greater distance from the nucleus and increased shielding.

Transition Metals

• Electrons added to d orbitals increase repulsion while adding protons increases attraction.
• Trends in transition metals are less predictable due to varying energy levels of d electrons.
• Overall trends usually remain intact despite individual element variations.