Compare electronegativity values across Pauling, Mulliken, and Allred-Rochow scales. Predict bond polarity, bond type, and dipole moments from electronegativity differences.
Electronegativity measures an atom's ability to attract bonding electrons toward itself. It is one of the most important periodic properties in chemistry, governing bond polarity, molecular shape, reactivity, and intermolecular forces. The Pauling scale (0.7 to 4.0) is the most widely used, but the Mulliken and Allred-Rochow scales offer complementary perspectives.
The electronegativity difference (ΔEN) between two bonded atoms predicts the bond type: ΔEN < 0.5 is nonpolar covalent, 0.5–1.7 is polar covalent, and > 1.7 is predominantly ionic. This simple rule lets you quickly assess any bond just by looking up two numbers on the periodic table.
This calculator compares electronegativity values for any two elements, predicts bond type and percent ionic character, shows periodic trends, and provides a reference table with Pauling values for common elements. It also converts between the Pauling, Mulliken, and Allred-Rochow scales.
For best results, combine calculator output with direct observation and periodic check-ins with a veterinarian or qualified advisor. Small adjustments made early usually improve comfort, safety, and long-term outcomes more than large corrective changes made later.
Quickly predict bond polarity and type from any pair of elements. Compare electronegativity scales and understand periodic trends. Essential for general and organic chemistry. This electronegativity calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
ΔEN = |EN_A - EN_B| Bond type: ΔEN < 0.5 → nonpolar covalent, 0.5-1.7 → polar covalent, >1.7 → ionic Percent ionic character ≈ [1 - e^(-0.25 × ΔEN²)] × 100% Mulliken: EN = (IE + EA) / 2 (in eV) Pauling: EN ∝ √(bond energy excess)
Result: ΔEN = 2.23, ionic bond (76% ionic character)
Sodium (EN = 0.93) and chlorine (EN = 3.16) have ΔEN = 3.16 - 0.93 = 2.23. This exceeds 1.7, predicting an ionic bond. Percent ionic character = [1 - e^(-0.25 × 2.23²)] × 100% = 76%.
Linus Pauling's scale (1932) derives electronegativity from bond dissociation energies. Robert Mulliken's scale (1934) uses the average of ionization energy and electron affinity. The Allred-Rochow scale (1958) estimates the electrostatic force exerted by the effective nuclear charge on bonding electrons. All three correlate well but differ in numerical values.
In organic chemistry, electronegativity differences determine functional group reactivity. Carbonyl groups (C=O) are reactive because oxygen pulls electron density away from carbon. Electronegativity also explains why fluorine substitution changes drug properties, why amines are basic, and why carbon-fluorine bonds are among the strongest single bonds known.
Group electronegativity assigns EN values to functional groups (e.g., -CF₃ is more electronegative than -CH₃). Electronegativity equalization models predict charge distribution in molecules. Allen's spectroscopic electronegativity uses average orbital energies for a more rigorous definition.
Fluorine, with a Pauling electronegativity of 3.98 — the highest of any element. It attracts bonding electrons more strongly than any other atom.
Francium or cesium at ~0.7 on the Pauling scale. Low electronegativity elements easily lose electrons and form cations.
No — it's an approximation. Many bonds near ΔEN = 1.7 show mixed character. Percent ionic character provides a more nuanced picture. Even NaCl has only ~76% ionic character.
Noble gases rarely form bonds, so electronegativity is typically not defined for them. Some scales assign values for Xe and Kr based on their known compounds.
The more electronegative atom in a bond carries a partial negative charge (δ-), the less electronegative gets δ+. This creates a bond dipole that contributes to molecular polarity.
As you move right across a period, nuclear charge increases while shielding stays roughly constant, so the effective nuclear charge on bonding electrons increases — atoms pull harder on shared electrons. This keeps planning practical and lowers the chance of preventable errors.