Learn & Review: Periodic Trends: Electronegativity, Ionization Energy, Atomic Radius

Summary of Periodic Trends

This summary outlines key periodic trends in chemistry, including electronegativity, ionization energy, electron affinity, atomic radii, ion size, and metallic character.

1. Electronegativity

• Definition: The ability of an element to attract electrons.
• Trend: Increases as you move across the periodic table (left to right) and up a column.
• Key Points:
  • Fluorine is the most electronegative atom.
  • Halogens generally have the highest electronegativity.
  • Noble gases are not electronegative because they have a full valence shell (octet) and are stable.
  • Elements with seven valence electrons (halogens) are highly electronegative as they are close to achieving a stable octet.
• Tiebreaker: If two elements are equidistant from fluorine, the one higher up on the periodic table is more electronegative.

2. Ionization Energy (IE)

• Definition: The energy required to remove an electron from an atom in its gaseous state.
• Trend: Similar to electronegativity, it increases across a period and up a group.
• Key Points:
  • This typically refers to the first ionization energy, the energy to remove the first electron.
  • Removing an electron from a neutral atom results in a positively charged ion (cation).
  • Electrons are removed one at a time.
• Examples:
  • Sulfur has a higher ionization energy than aluminum because sulfur is closer to fluorine.
  • Nitrogen has a higher ionization energy than silicon because nitrogen is higher up and closer to fluorine.
• Exceptions:
  • Equidistant Elements: If elements are the same distance from fluorine, ionization energy cannot be determined solely by this trend.
  • Group 2A to 3A (e.g., Beryllium to Boron): Elements in Group 2A (like Beryllium) have higher ionization energies than those in Group 3A (like Boron) despite the general trend. This is due to the stability of a full s orbital in Group 2A elements.
  • Group 5A to 6A (e.g., Nitrogen to Oxygen): Elements in Group 5A (like Nitrogen) have higher ionization energies than those in Group 6A (like Oxygen). This is due to the stability of a half-filled p orbital in Group 5A elements.

3. Electron Affinity

• Definition: The energy change that occurs when an electron is added to a neutral atom in its gaseous state. It's essentially the opposite of ionization energy.
• Trend: Generally increases as you move across the periodic table (left to right) and up a column, meaning elements closer to fluorine have a higher electron affinity.
• Key Points:
  • Adding an electron to a neutral atom results in a negatively charged ion (anion).
  • A more negative electron affinity value indicates a higher affinity for electrons.
• Exceptions:
  • Group 1A to 2A: Elements in Group 1A (like Lithium) have higher electron affinities than those in Group 2A (like Beryllium).
  • Group 4A to 5A: Elements in Group 4A (like Carbon) have higher electron affinities than those in Group 5A (like Nitrogen).
  • Noble gases have low electron affinities.
  • Halogens have the highest electron affinities.

4. Atomic Radii

• Definition: Half the distance between the nuclei of two identical atoms bonded together.
• Trend: The trend is the opposite of electronegativity, ionization energy, and electron affinity. Atomic radii increase as you move from right to left across the periodic table and down a column.
• Key Points:
  • Elements in the bottom-left corner of the periodic table (like Cesium, Francium) have the largest atomic radii.
• Examples:
  • Nitrogen has a larger atomic radii than Fluorine.
  • Aluminum has a larger atomic radii than Nitrogen.

5. Ion Size

• Definition: The size of an atom when it has gained or lost electrons to form an ion.
• General Trend: Increases as you move down a column.
• Key Factors:
  • Anions (Negative Charge): Generally larger than their neutral atom counterparts because adding electrons increases electron-electron repulsion, expanding the electron cloud.
  • Cations (Positive Charge): Generally smaller than their neutral atom counterparts because losing electrons reduces electron-electron repulsion and the remaining electrons are pulled closer to the nucleus.
• Isoelectronic Species: Elements with the same number of electrons and thus the same electron configuration.
• Size Order: Anions > Neutral Atoms > Cations.
• Tiebreaker for Ions with Same Charge: If ions have the same charge, the one with more electrons will be larger.
• Tiebreaker for Ions with Different Charges: The charge is the primary factor. A higher negative charge (more negative anion) leads to a larger size, while a higher positive charge (more positive cation) leads to a smaller size.

6. Metallic Character

• Definition: How closely an element's properties match those of a metal.
• Properties of Metals: Malleable, good conductors of electricity, shiny, reflect light, ionize easily.
• Trend: Increases as you move from right to left across the periodic table and down a column. This trend is similar to atomic radii.
• Key Points:
  • Nonmetals (on the right side of the periodic table) have low metallic character.
  • Metals (on the left and center of the periodic table) have high metallic character.
• Examples:
  • An element lower in a column will have more metallic character than one above it.
  • An element further to the left on the periodic table will have more metallic character than one to its right.