Find the electronegativity of any element on the Pauling scale. Select two elements to calculate the electronegativity difference, bond polarity, percentage ionic character, and dipole moment direction.
Electronegativity is the tendency of an atom to attract the shared electron pair in a chemical bond toward itself. Linus Pauling defined the scale in 1932 based on bond dissociation energies: if a bond A–B releases more energy than would be expected from purely covalent A–A and B–B bonds, that extra energy reflects an ionic resonance stabilization caused by the electronegativity difference.
The electronegativity difference (ΔEN) between bonded atoms determines bond character. At ΔEN = 0, the bond is purely nonpolar covalent. As ΔEN increases, the bond becomes increasingly polar — electrons are pulled toward the more electronegative atom. At approximately ΔEN ≥ 1.7, the bond has more than 50% ionic character and is classified as ionic.
Bond Type Classification by ΔEN
ΔEN Range
Bond Type
Example
% Ionic
0.0
Nonpolar covalent
H–H, Cl–Cl
0%
0.1–0.4
Slightly polar covalent
C–H (0.35)
<5%
0.5–1.6
Polar covalent
H–Cl (0.96), H–O (1.24)
5–50%
≥1.7
Ionic
Na–Cl (2.23)
>50%
Electronegativity Trends in the Periodic Table
Electronegativity increases across a period because nuclear charge increases while atomic radius decreases — bonding electrons experience a stronger pull toward the nucleus. It decreases down a group because atomic radius increases, placing bonding electrons farther from the increasingly shielded nuclear charge. Fluorine at the top-right corner of the periodic table has the highest EN (3.98). Cesium and francium at the bottom-left have the lowest (<0.8).
💡 Most Electronegative Elements (Pauling scale):
F = 3.98 › O = 3.44 › Cl = 3.16 › N = 3.04 › Br = 2.96 › I = 2.66 › S = 2.58 › C = 2.55 › H = 2.20
Noble gases: traditionally listed as 0 or not assigned (no stable bonds formed)
Frequently Asked Questions
Tendency of an atom to attract bonding electrons. Pauling scale: 0.7 (Fr) to 3.98 (F). Higher EN = stronger electron attraction in bonds.
Based on bond dissociation energies (1932). Dimensionless, ranges 0.7–3.98. Most widely used EN scale. Fluorine set as 3.98 as reference.
ΔEN=0: nonpolar covalent. ΔEN 0.1–1.6: polar covalent. ΔEN ≥1.7: ionic (>50% ionic character). Boundary is approximate — ionic/covalent is a continuum.
% ionic = 16|ΔEN| + 3.5(ΔEN)² (Hanney-Smyth). At ΔEN=1.7: ~50% ionic. NaCl ΔEN=2.23: ~70% ionic. Even NaCl has partial covalent character.
Fluorine (F) = 3.98 (Pauling). Extremely small radius + high nuclear charge = strongest electron attraction. Order: F > O > Cl > N > Br > I > S > C.
Increases left→right (higher Z-eff, smaller radius). Decreases top→bottom (larger radius, electrons farther from nucleus). F (top-right) highest; Cs/Fr (bottom-left) lowest.
H = 2.20 (Pauling). Close to C (2.55), so C–H bonds are nearly nonpolar. H is more EN than most metals but less than nonmetals, allowing it to bond as H⁺ or H⁻.
Pauling: based on bond energies, dimensionless. Mulliken: EN = (IE + EA)/2, in eV. Allred-Rochow: based on Z-eff and covalent radius. All show similar trends; Pauling is most widely used.
Unequal electron sharing due to ΔEN. More EN atom gets δ⁻, less EN atom gets δ⁺. Examples: H–Cl (ΔEN=0.96), H–O (ΔEN=1.24), H–N (ΔEN=0.84).
Predicts: bond polarity and dipole direction; bond type (ionic/covalent); oxidation states; acid-base strength; reactivity patterns; molecular polarity from geometry.