Periodic Trends Lecture Notes

Overview of Periodic Trends

• Atomic Radius
• Ionic Radius
• Electronegativity
• Ionization Energy
• Electron Affinity
• Metallic Character

Atomic Radius

General Trends

• Atomic size increases:
  • Left to Right: Decreases
  • Top to Bottom: Increases

Example: Hydrogen vs. Helium

• Hydrogen (H): Larger atomic radius than Helium (He) despite lower atomic mass.
  • H: Atomic mass = 1, He: Atomic mass = 4
• Reason:
  • More protons in He create stronger electrostatic attraction, pulling electrons closer to the nucleus, reducing atomic size.

Comparing Hydrogen and Lithium

• Lithium (Li): Has more protons than H but is larger due to having additional electron shells.
• Key Concept:
  • Size increases down a group due to additional shells of electrons.

Core Electrons and Shielding

• Inner core electrons shield outer electrons from nuclear charge, affecting atomic size.

Ranking Atomic Sizes

Example ranking of elements:

• Chlorine (Cl), Magnesium (Mg), Phosphorus (P):
  • Rank: Mg > P > Cl

Ionic Radii

Cations and Anions

• Cations (e.g., Li+): Smaller than their parent atoms due to loss of an electron and electron shell.
• Anions (e.g., N3-): Larger than their parent atoms due to added electrons causing electron repulsion.

Trends

• Ionic size increases: Similar trends as atomic size.
• Isoelectronic ions:
  • Compare size based on nuclear charge; more protons = smaller size.
• Example ranking:
  • Na+ < Mg2+ < Al3+ (smaller with more positive charge)

Electronegativity

Definition

• Ability of an atom to attract electrons.

General Trend

• Increases towards Flourine (F):
  • Upward and Rightward: Increases
• Metals are electropositive (tend to lose electrons), while nonmetals are electronegative (tend to gain electrons).

Example Ranking

• Elements: Silicon (Si), Magnesium (Mg), Chlorine (Cl), Aluminum (Al):
  • Rank: Mg < Al < Si < Cl

Ionization Energy

Definition

• Energy required to remove an electron from a gaseous atom.

Trends

• Increases towards Helium (He):
  • Upward and Rightward: Increases
• Metals have lower ionization energy than nonmetals.

Example: Lithium vs. Beryllium

• Beryllium (Be) has higher ionization energy than Lithium (Li) due to higher nuclear charge and smaller size.

Key Takeaways

• Distance between electrons and nucleus affects ionization energy significantly.

Electron Affinity

Definition

• Energy change when adding an electron to a gaseous atom.
  • Generally exothermic for nonmetals (e.g., halogens).

Trends

• Increases (more exothermic) towards Group 7 (halogens):
  • Group 1: Less exothermic
  • Group 2 and Group 8: Often endothermic

Metallic Character

General Trends

• Increases towards Bottom Left:
  • Metals exhibit higher metallic character than nonmetals.

Example Ranking

• Elements: Silicon (Si), Sodium (Na), Sulfur (S), Aluminum (Al), Chlorine (Cl):
  • Rank: Cl < S < Si < Al < Na

Summary of Key Points

• Atomic Radius: Increases down a group, decreases across a period.
• Ionic Radius: Cations smaller than parent; anions larger.
• Electronegativity: Increases towards F, nonmetals > metals.
• Ionization Energy: Increases upward and rightward; influenced by nuclear charge and distance.
• Electron Affinity: Generally exothermic for nonmetals, varies by group.
• Metallic Character: Increases down and to the left in the periodic table.