Electronegativity & Bond Polarity

Electronegativity measures how strongly an atom attracts electrons in a covalent bond. The Pauling scale (1932) is the most widely used, ranging from Cs and Fr at 0.79 to F at 3.98.

Fluorine is the most electronegative element. The scale is dimensionless and based on bond dissociation energies. Noble gases have undefined or zero electronegativity because they rarely form bonds.

Electronegativity generally increases across a period (left to right) and decreases down a group in the periodic table.

The electronegativity difference (Δχ) between two bonded atoms determines the bond type.

• Nonpolar covalent: $\Delta\chi < 0.5$ — equal sharing (H-H, C-C)
• Polar covalent: $0.5 \le \Delta\chi \le 1.8$ — unequal sharing (H-F, O-H, C-O)
• Ionic: $\Delta\chi > 1.8$ — electron transfer (Na-Cl, K-F)

These thresholds are guidelines, not sharp rules. Real bonds exist on a continuum from purely covalent to fully ionic.

A bond dipole moment μ is a vector pointing from the partial positive charge toward the partial negative charge. It is measured in Debye (D), where 1 D = 3.336 × 10^-30 C·m.

where $q$ is the partial charge (in coulombs) and $d$ is the bond length (in meters). For a rough estimate, $q \approx \delta \cdot e$ using the partial charge fraction δ.

Molecular geometry determines whether individual bond dipoles cancel or add. CO₂ has two polar C=O bonds but zero net dipole because they are antiparallel.

Percent ionic character estimates what fraction of a bond resembles a full ionic bond. The Hannay-Smyth equation uses the electronegativity difference:

For H-Cl (Δχ = 0.96), percent ionic character is about 19%. For Na-Cl (Δχ = 2.23), it exceeds 70%, reflecting near-complete electron transfer.

The partial charge δ on each atom equals the percent ionic character divided by 100. A value of 0.5 means each atom carries half an elementary charge.