How does atomic radius vary across a period?

Short Answer

Across a period, the atomic radius decreases from left to right in the periodic table. This is because the nuclear charge increases as the number of protons increases, pulling the electrons closer to the nucleus.

Even though electrons are being added to the same energy level, the increasing attraction between the nucleus and electrons dominates, causing the atom to shrink. For example, in Period 2, lithium has a larger atomic radius than fluorine.

Detailed Explanation :

Trend of Atomic Radius Across a Period

The atomic radius is the distance from the nucleus to the outermost electron. When moving across a period from left to right:

  1. Number of Protons Increases:
    • The atomic number increases across a period, meaning more protons in the nucleus.
    • This increases the nuclear charge, which pulls electrons closer to the nucleus.
  2. Electrons Are Added to Same Shell:
    • Additional electrons enter the same principal energy level (same shell).
    • These electrons do not significantly increase electron shielding.
  3. Effective Nuclear Charge Increases:
    • The pull of the nucleus on the valence electrons becomes stronger.
    • This reduces the size of the electron cloud, shrinking the atomic radius.

Explanation Using Examples

  • Period 2 Elements: Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne)
    • Lithium → largest radius in the period
    • Fluorine → smallest radius in the period
  • This shows a gradual decrease across the period.

Factors Affecting Atomic Radius Across a Period

  1. Nuclear Charge (Z):
    • Increasing number of protons → stronger pull → smaller radius.
  2. Electron Shielding:
    • Inner electrons shield outer electrons partially, but since electrons are added to the same shell, shielding does not increase enough to offset nuclear pull.
  3. Electron-Electron Repulsion:
    • Slight repulsion between added electrons occurs, but the effect is smaller than the nuclear attraction.
  4. Number of Electron Shells:
    • Remains constant across a period → no increase in size due to shells.

Trends in Chemical and Physical Properties

  1. Ionization Energy:
    • As atomic radius decreases, electrons are closer to the nucleus → more energy required to remove an electron.
    • Example: Lithium has lower ionization energy than fluorine.
  2. Electronegativity:
    • Smaller radius → stronger pull on bonding electrons → higher electronegativity across a period.
  3. Metallic Character:
    • Metals (on the left) have larger radii and lose electrons easily → more metallic.
    • Non-metals (on the right) have smaller radii → less metallic.

Significance of Periodic Trend

  • Helps predict chemical reactivity: Smaller atoms like fluorine are highly reactive non-metals.
  • Explains bond length in molecules: Smaller atoms form shorter bonds.
  • Supports understanding of electron configuration and periodic properties.

Exceptions

  • Minor anomalies may occur due to electron repulsion in half-filled or full-filled orbitals, but overall trend of decreasing atomic radius across a period is consistent.
Conclusion

Across a period in the periodic table, atomic radius decreases from left to right due to the increasing nuclear charge pulling electrons closer while electrons are added to the same energy level. This trend influences ionization energy, electronegativity, metallic character, and chemical reactivity, making it an essential concept in understanding periodic properties of elements.