Trends in the Periodic Table

In this lecture, we explore the various trends in the periodic table. A trend is a pattern that can be observed across the periodic table either from top to bottom within a group or from left to right within a period. We focused on three key properties: atomic radius, ionization energy, and electronegativity.

Atomic Radius

• Definition: The atomic radius refers to the size of an atom, specifically the distance from the nucleus to the outermost electron in the electron cloud.
• Trend within a Group: From top to bottom, the atomic radius increases.
  • This increase occurs because more electron shells are required to accommodate the increasing number of electrons.
  • Analogy: Think of a snowman, which grows larger with more layers.
• Trend within a Period: From left to right, the atomic radius decreases.
  • Despite adding electrons, additional protons increase nuclear attraction, pulling electrons closer to the nucleus.
  • Example: In period two, lithium has the largest atomic radius, and neon has the smallest.

Ionization Energy

• Definition: Ionization energy is the energy required to remove an electron from an atom.
• Trend across a Period: Ionization energy increases from left to right.
  • Smaller atoms (e.g., neon) hold electrons more tightly, requiring more energy to remove an electron.
• Trend within a Group: Ionization energy decreases from top to bottom.
  • Larger atoms with more electron shielding make it easier to remove an electron.
  • Shielding: Inner electrons block the nucleus’s pull, making valence electrons easier to remove.

Electronegativity

• Definition: Electronegativity measures how strongly an atom attracts electrons.
• Scale: Ranges from 0 (does not want electrons) to 4 (highly attracts electrons).
• Trend across a Period: Electronegativity increases from left to right.
  • Elements like fluorine, which need one more electron to complete their valence shell, have high electronegativity.
  • Example: Fluorine (3.98) vs. Lithium (0.98).
• Trend within a Group: Electronegativity decreases from top to bottom.
  • Bigger atoms have a weaker pull on electrons due to increased distance between nucleus and valence electrons.

Summary

Understanding the reason behind these trends is crucial, often relating back to atomic size and the distance between the nucleus and valence electrons. These trends are fundamental in predicting and understanding chemical behavior.