Periodic trends in atomic properties reveal fascinating patterns across the periodic table. These trends, including atomic size, ionization energy, and electron affinity, are influenced by factors like nuclear charge and electron shielding.

Understanding these trends helps predict element behavior and reactivity. By comparing atomic and ionic sizes, we gain insights into how elements form bonds and interact in chemical reactions. This knowledge is crucial for understanding chemical properties and reactions.

Periodic Trends in Atomic Properties

Periodic trends in atomic size, ionization energy, and electron affinity across periods and down groups of the periodic table

• Atomic size (atomic radius)
  • Decreases across a period from left to right due to increasing effective nuclear charge with same number of electron shells (Li to Ne)
  • Increases down a group due to increasing number of electron shells and shielding effect of inner electrons (Li to Cs)
• Ionization energy
  • Energy required to remove an electron from a gaseous atom
  • Increases across a period from left to right due to increasing effective nuclear charge causing electrons to be more strongly attracted to nucleus (Li to Ne)
  • Decreases down a group due to increasing atomic size and valence electrons being farther from nucleus (Li to Cs)
• Electron affinity
  • Energy change when an electron is added to a gaseous atom
  • Becomes more negative across a period from left to right due to increasing effective nuclear charge causing electrons to be more strongly attracted to nucleus (F to Cl)
  • Becomes less negative down a group due to increasing atomic size and valence electrons being farther from nucleus (F to I)
• Electronegativity
  • Follows similar trends to ionization energy and electron affinity
  • Increases across a period and decreases down a group

Atomic vs ionic size comparisons

• Cations (positive ions)
  • Smaller than their parent atoms due to electrons being removed from the outermost shell, resulting in increased effective nuclear charge (Na vs Na+)
  • Higher positive charge leads to smaller ionic radius due to more electrons being removed and greater effective nuclear charge (Mg2+ vs Al3+)
• Anions (negative ions)
  • Larger than their parent atoms due to electrons being added to the outermost shell, resulting in decreased effective nuclear charge (Cl vs Cl-)
  • Higher negative charge leads to larger ionic radius due to more electrons being added and lesser effective nuclear charge (O2- vs S2-)

Nuclear charge and periodic trends

• Effective nuclear charge ($Z_{eff}$)
  • Net positive charge experienced by an electron
  • Increases across a period due to same number of electron shells and increasing number of protons (Li to Ne)
  • Decreases down a group due to increasing number of electron shells and shielding effect of inner electrons (Li to Cs)
• Influence on periodic trends
  • Atomic size: Higher $Z_{eff}$ leads to smaller atomic size as electrons are more strongly attracted to nucleus (F vs Cl)
  • Ionization energy: Higher $Z_{eff}$ leads to higher ionization energy as electrons are more difficult to remove (N vs P)
  • Electron affinity: Higher $Z_{eff}$ leads to more negative electron affinity as electrons are more strongly attracted to nucleus (O vs S)
  • Ionic size: Higher $Z_{eff}$ leads to smaller cations (Na+ vs K+) and larger anions (F- vs Cl-) due to increased or decreased attraction between electrons and nucleus

Periodic Table Organization and Element Properties

• The periodic table is organized based on atomic number, which increases from left to right and top to bottom
• Electron configuration determines an element's position in the periodic table and influences its properties
• Metallic character generally increases down a group and decreases across a period from left to right