How does electronegativity vary in groups and periods?

Short Answer
Electronegativity varies predictably in the periodic table.
  • Across a period (left to right): Electronegativity increases because nuclear charge increases while atomic radius decreases, so atoms attract electrons more strongly.
  • Down a group (top to bottom): Electronegativity decreases because additional electron shells increase atomic size and shielding, reducing the nucleus’s pull on bonding electrons.

These trends explain bond polarity, chemical reactivity, and periodic properties of elements.

Detailed Explanation :

Definition of Electronegativity

Electronegativity (EN) is the ability of an atom to attract shared electrons in a chemical bond. It is a dimensionless quantity and can be measured using the Pauling scale.

  • High EN → strong pull on electrons.
  • Low EN → weak pull on electrons.
  • Understanding electronegativity trends helps explain bond type, molecular polarity, and reactivity.

Variation Across a Period

  1. Increase from Left to Right:
    • As we move across a period, atomic number increases, adding protons to the nucleus.
    • Electrons are added to the same shell → shielding remains nearly constant.
    • The effective nuclear charge increases, pulling bonding electrons closer.
  2. Effect on Atomic Behavior:
    • Atoms on the left (alkali metals) → low EN → weakly attract electrons.
    • Atoms on the right (halogens) → high EN → strongly attract electrons.
  3. Examples:
    • Period 2: Lithium (EN 0.98) → Fluorine (EN 3.98)
    • Shows steady increase in electronegativity across the period.
  4. Reasoning:
    • Smaller atomic radius → electrons closer to nucleus → higher EN.
    • Increased nuclear charge → stronger attraction for electrons.

Variation Down a Group

  1. Decrease from Top to Bottom:
    • As we move down a group, new electron shells are added, increasing the distance between nucleus and bonding electrons.
    • Inner electrons shield outer electrons from nuclear attraction.
    • Result: The ability to attract electrons decreases → lower electronegativity.
  2. Effect on Atomic Behavior:
    • Elements at the top of a group (e.g., fluorine) → high EN → strongly attracts electrons.
    • Elements at the bottom of a group (e.g., iodine) → lower EN → weaker attraction for electrons.
  3. Examples:
    • Group 17: Fluorine (EN 3.98), Chlorine (EN 3.16), Bromine (EN 2.96), Iodine (EN 2.66)
    • Shows gradual decrease down the group.
  4. Reasoning:
    • Larger atomic size → bonding electrons farther from nucleus.
    • Increased shielding → reduced nuclear pull → lower EN.

Factors Influencing Variation

  1. Atomic Size: Smaller atoms → higher EN; larger atoms → lower EN.
  2. Nuclear Charge: Higher charge → higher EN across periods.
  3. Electron Shielding: More inner electrons → lower EN down groups.
  4. Electron Configuration: Half-filled and filled orbitals resist electron gain → lower EN.

Significance of Trends

  1. Predicting Bond Polarity:
    • Large EN difference → ionic bond.
    • Small EN difference → polar covalent bond.
    • No difference → non-polar covalent bond.
  2. Chemical Reactivity:
    • High EN → atoms attract electrons → strong oxidizing agents.
    • Low EN → atoms lose electrons easily → strong reducing agents.
  3. Periodic Behavior:
    • Explains reactivity trends, metallic/non-metallic character, and molecular properties.
  4. Industrial Applications:
    • Useful in chemical synthesis, material design, and drug development.
Conclusion

Electronegativity increases across a period due to higher nuclear charge and smaller atomic size, and decreases down a group due to added electron shells and shielding. These trends are essential for predicting bond type, polarity, chemical reactivity, and periodic properties of elements, making electronegativity a key concept in chemistry.