Question

1. Describe electronegativity, and explain the trend of electronegativity in the Periodic Table.

Answer 1

Electronegativity is a measure of an atom’s tendency to attract and hold bonding electrons when it is chemically bonded to another atom. It is a relative, dimensionless quantity (no units) most commonly reported on the Pauling scale.

Why electronegativity changes across the Periodic Table
- Across a period (left → right): electronegativity increases.
- Reason: nuclear charge (number of protons) increases while valence electrons occupy roughly the same principal energy level. The increased effective nuclear charge pulls bonding electrons more strongly, and atomic radius decreases, so atoms attract shared electrons more strongly.
- Down a group (top → bottom): electronegativity decreases.
- Reason: valence electrons occupy higher principal energy levels that are farther from the nucleus and experience greater shielding by inner electrons. The larger atomic radius and increased shielding reduce the nucleus’s ability to attract bonding electrons.

Additional points
- The most electronegative element is fluorine (Pauling ≈ 3.98); alkali metals (e.g., francium, caesium) are among the least electronegative.
- Electronegativity correlates with ionization energy and electron affinity: atoms with high ionization energies and high (more positive) electron affinities tend to be more electronegative.
- Electronegativity differences predict bond character: large differences → ionic character; small differences → covalent (polar or nonpolar) character.
- Noble gases are often excluded or treated separately because they rarely form bonds (exceptions like xenon compounds exist).
- Other scales (Mulliken, Allred–Rochow) exist, but trends are the same.

Example: In a period, sodium (Na) is much less electronegative than chlorine (Cl), so Cl attracts bonding electrons strongly and Na–Cl bonding is highly polar/ionic.