Periodic Table Trends to Know for AP Chemistry

Understanding periodic table trends is essential in AP Chemistry and Inorganic Chemistry I. These trends, like atomic and ionic radii, electronegativity, and ionization energy, help explain how elements interact, bond, and behave in chemical reactions.

1. Atomic radius

   • Increases down a group due to the addition of electron shells.
   • Decreases across a period from left to right as effective nuclear charge increases.
   • Influences the size of atoms and their ability to form bonds.
   • Larger atomic radius generally leads to lower ionization energy.

2. Ionic radius

   • Cations (positively charged ions) are smaller than their neutral atoms due to loss of electron shells.
   • Anions (negatively charged ions) are larger than their neutral atoms due to the addition of electrons.
   • Trends in ionic radius follow similar patterns to atomic radius across periods and groups.
   • Important for understanding ionic bonding and lattice structures.

3. Electronegativity

   • Measures an atom's ability to attract electrons in a bond.
   • Increases across a period due to higher effective nuclear charge.
   • Decreases down a group as increased distance from the nucleus reduces attraction.
   • Key factor in determining bond polarity and molecular behavior.

4. Electron affinity

   • Refers to the energy change when an electron is added to a neutral atom.
   • Generally becomes more negative (more favorable) across a period.
   • Less predictable down a group; noble gases have positive electron affinities.
   • Important for understanding an element's tendency to form anions.

5. Ionization energy

   • The energy required to remove an electron from a neutral atom.
   • Increases across a period due to increased nuclear charge and decreased atomic radius.
   • Decreases down a group as electrons are further from the nucleus and experience more shielding.
   • Critical for predicting reactivity and the formation of cations.

6. Metallic character

   • Increases down a group as atomic size increases and ionization energy decreases.
   • Decreases across a period from left to right as electronegativity increases.
   • Metals tend to lose electrons easily, while nonmetals tend to gain them.
   • Influences the physical and chemical properties of elements.

7. Reactivity

   • For metals, reactivity increases down a group and decreases across a period.
   • For nonmetals, reactivity generally increases across a period and decreases down a group.
   • Influenced by ionization energy and electronegativity.
   • Important for predicting chemical reactions and compound formation.

8. Melting and boiling points

   • Vary widely across the periodic table; metals typically have higher melting and boiling points.
   • Generally increase across a period due to stronger metallic bonds and covalent network structures.
   • Decrease down a group for nonmetals due to weaker van der Waals forces.
   • Important for understanding the physical state of elements at room temperature.

9. Electron configuration

   • Describes the distribution of electrons in an atom's orbitals.
   • Follows the Aufbau principle, Hund's rule, and Pauli exclusion principle.
   • Determines an element's chemical properties and reactivity.
   • Important for predicting bonding behavior and the formation of ions.

10. Effective nuclear charge

    • The net positive charge experienced by an electron in a multi-electron atom.
    • Increases across a period, leading to a stronger attraction between the nucleus and electrons.
    • Decreases down a group due to increased shielding from inner electrons.
    • Key concept for understanding trends in atomic and ionic radii, ionization energy, and electronegativity.