Overview
This lecture covers periodic properties, focusing on atomic and ionic radius, how these radii are measured, their trends in the periodic table, and key exceptions to the trends.
Periodic Properties and Atomic Radius
- Periodic properties repeat at regular intervals in elements.
- Atomic radius is the distance from the nucleus to the outermost electron shell; not well-defined due to the wavelike nature of electrons.
- Modern measurement defines atomic radius in bonded states (not isolated atoms).
Types of Atomic Radii
- Covalent Radius: Half the distance between nuclei of two identical atoms joined by a covalent bond.
- For different atoms: bond length = RA + RB (if electronegativity is the same).
- If electronegativity differs: bond length = RA + RB – 0.9(χA – χB).
- Van der Waals Radius: Half the closest distance between nuclei of two non-bonded atoms/molecules; used for noble gases.
- Van der Waals radius is always greater than covalent radius.
- Metallic Radius: Half the distance between nuclei of adjacent atoms in a metallic lattice.
- Size order: van der Waals > metallic > covalent.
Trends in Atomic Radius
- Across a period (left to right): atomic radius decreases due to increasing nuclear charge (stronger attraction pulls electrons closer).
- Down a group: atomic radius increases due to added electron shells increasing distance from nucleus.
Exceptions in Atomic Radius Trends
- Noble gases have larger atomic radii (van der Waals) compared to adjacent elements (covalent radius).
- In D-block (transition metals), atomic radius decreases at first, then becomes almost constant, then slightly increases due to electron-electron repulsion balancing nuclear charge.
- Lanthanoid Contraction: In F-block elements, poor shielding by 4f electrons causes a sharp decrease in atomic radius, so 4d and 5d elements have similar radii.
Shielding/Screening Effect
- Inner electrons shield outer electrons from the nucleus, reducing effective nuclear charge felt by outer electrons.
- Shielding effectiveness: s > p > d > f (f is poorest at shielding).
Ionic Radius
- Cations (positive ions) are smaller than their parent atom due to increased nuclear pull on fewer electrons.
- Anions (negative ions) are larger than their atom due to greater electron-electron repulsion and reduced nuclear pull per electron.
- Size order: anion > atom > cation.
- Higher positive charge = smaller cation; higher negative charge = larger anion.
Trends and Isoelectronic Series
- For same charge, ionic radius decreases across a period (increasing nuclear charge).
- Down a group, ionic radius increases (more shells).
- Isoelectronic species (same electron count): Larger atomic number means smaller ionic radius (size ∝ 1/atomic number).
Key Terms & Definitions
- Periodic Property — A property that repeats its trend at regular intervals in the periodic table.
- Atomic Radius — Distance from the nucleus to the outermost electron shell.
- Covalent Radius — Half the distance between nuclei of two identical atoms bonded together.
- Van der Waals Radius — Half the distance between nuclei of two adjacent non-bonded atoms.
- Metallic Radius — Half the distance between nuclei of adjacent atoms in a metallic crystal.
- Ionic Radius — Effective radius of an ion in a crystal structure.
- Shielding Effect — The reduction of nuclear attraction on outer electrons due to inner electrons.
- Lanthanoid Contraction — Significant decrease in atomic radius across the lanthanide series due to poor f-orbital shielding.
- Isoelectronic Species — Ions or atoms with the same number of electrons.
Action Items / Next Steps
- Review and memorize the formulas for calculating covalent radius.
- Practice arranging elements/ions based on atomic and ionic radii, including exceptions and isoelectronic species.
- Study the concept and implications of lanthanoid contraction and shielding effect.