Last updated: June 2026
Use this bond polarity and electronegativity calculator to compare two atoms, calculate their electronegativity difference and estimate whether the bond is approximately non-polar covalent, polar covalent or strongly ionic in character.
What is electronegativity?
Electronegativity describes how strongly an atom attracts the shared pair of electrons in a covalent bond. On the Pauling scale, fluorine has the highest electronegativity.
When two bonded atoms have different electronegativities, the bonding electrons spend more time closer to the more electronegative atom. This produces a bond dipole and partial charges.
How bond polarity is calculated
The electronegativity difference is found using:
Δχ = |χ1 − χ2|
The absolute value is used because electronegativity difference cannot be negative. The calculator then compares the result with broad school-level guidelines.
| Electronegativity difference | Approximate interpretation |
|---|---|
| 0.0 to about 0.4 | Non-polar or only weakly polar covalent |
| Above about 0.4 to 1.7 | Polar covalent |
| Above about 1.7 | Large ionic character; often ionic for a metal–non-metal pair |
These ranges are deliberately described as approximate. A bond does not suddenly change character at one exact electronegativity difference.
How to assign δ+ and δ−
The more electronegative atom attracts the bonding electrons more strongly and is labelled δ−. The other atom is electron-deficient relative to it and is labelled δ+.
For H–Cl, hydrogen has electronegativity 2.20 and chlorine has electronegativity 3.16. Therefore:
Δχ = |3.16 − 2.20| = 0.96
H–Cl is polar covalent, with Hδ+ and Clδ−.
What is percent ionic character?
Percent ionic character can be estimated from electronegativity difference using the empirical expression:
percent ionic character = [1 − e−0.25(Δχ)2] × 100%
The value is an estimate of how strongly charge separation contributes to the bond. It should not be treated as a strict division between covalent and ionic bonding.
Worked examples
C–H bond
Carbon has electronegativity 2.55 and hydrogen has electronegativity 2.20, so Δχ = 0.35. The C–H bond is commonly treated as almost non-polar in introductory chemistry.
H–Cl bond
Chlorine has electronegativity 3.16 and hydrogen has electronegativity 2.20, so Δχ = 0.96. The bond is polar covalent, with chlorine carrying δ−.
Na–Cl interaction
Sodium has electronegativity 0.93 and chlorine has electronegativity 3.16, so Δχ = 2.23. Because this is also a metal–non-metal pair, school chemistry treats sodium chloride as ionic.
Common exam traps
Frequently asked questions
Does a larger electronegativity difference always mean a stronger bond?
No. Electronegativity difference describes the distribution of bonding electrons, not bond strength by itself. Bond strength also depends on bond order, atomic size and orbital overlap.
Can a molecule with polar bonds be non-polar?
Yes. In a symmetrical molecule, equal bond dipoles can cancel. Carbon dioxide is the standard example: each C=O bond is polar, but the linear molecule is non-polar overall.
Why are noble gases not included?
Noble-gas electronegativities are often omitted from introductory Pauling-scale tables because they do not commonly form the simple bonds considered at school level.
Is percent ionic character exact?
No. It is calculated from an empirical relationship and is best used as a comparison rather than an exact physical description of a bond.