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Bond Polarity and Dipole Checker

In short: Pick two elements to get the Pauling electronegativity difference, the bond type (nonpolar, polar covalent or ionic), which atom is δ− and the % ionic character. Or enter a molecule to see if it is polar, from its VSEPR shape and the vector sum of its bond dipoles. It is free and works in your web browser.

Pick two elements to see the electronegativity difference, bond type, which end is δ+ and δ−, and the % ionic character. Or enter a molecule (formula such as XeF4, or a SMILES string) to find out whether it is polar, with the VSEPR shape, the bond dipoles and their vector sum drawn out.

Quick answer: ΔEN for H–Cl = 3.16 − 2.20 = 0.96, so the bond is polar covalent with Cl δ− and H δ+; % ionic = (1 − e^(−0.96²/4)) × 100 = 20.6%. Water is bent with two lone pairs on O, so its O–H dipoles add up: H₂O is polar (1.85 D). CO₂ is linear, so its dipoles cancel: nonpolar.

Element AElement B

Examples:

Bond polarity: formulas and cut-offs

ΔEN = |χA − χB| % ionic character = (1 − e^(−ΔEN²/4)) × 100 (Pauling)

ΔEN (Pauling)	Bond type	Examples
less than 0.4	nonpolar covalent	Cl–Cl (0), C–H (0.35), C–S (0.03)
0.4 to 1.7	polar covalent	H–Cl (0.96), C–O (0.89), O–H (1.24)
1.7 to 2.0	borderline: ionic in most school books, polar covalent in some	H–F (1.78), B–F (1.94)
more than 2.0	ionic	Na–Cl (2.23), K–F (3.16)

This page uses the common convention: below 0.4 nonpolar, 0.4 to 1.7 polar covalent, above 1.7 ionic, and it flags 1.7 to 2.0 as borderline because some textbooks put the cut-off at 2.0. Bonds between two non-metals stay covalent even when ΔEN is large (H–F). The more electronegative atom is δ−, and the dipole arrow points toward it (the crossed tail marks the δ+ end).

Molecular polarity in three steps

Draw the Lewis structure and find the shape with VSEPR.
Draw a dipole arrow for each polar bond, pointing to the δ− atom.
Add the arrows as vectors. Symmetric shapes (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, square planar) with identical outer atoms cancel: nonpolar. Lone pairs on the centre or different outer atoms usually leave a net dipole: polar.

Worked examples

H–F bond

ΔEN = 3.98 − 2.20 = 1.78.

% ionic = (1 − e^(−0.79)) × 100 = 54.7%.

Polar covalent (borderline), F δ−

Water, H₂O

AX₂E₂, bent (104.5°). Both O–H dipoles (ΔEN 1.24) point toward O and partly the same way.

Polar, 1.85 D

BF₃

AX₃, trigonal planar. Three equal B–F dipoles at 120° add to zero.

Nonpolar (polar bonds, nonpolar molecule)

NH₃ vs NF₃

Both trigonal pyramidal. In NH₃ the lone pair and the N–H dipoles point the same way (1.47 D); in NF₃ they oppose (0.24 D).

Both polar

PCl₅ (NEET style)

AX₅, trigonal bipyramidal: the two axial dipoles cancel and the three equatorial ones cancel.

Nonpolar

CH₂Cl₂

Tetrahedral, but two C–Cl (0.61) and two C–H (0.35) bonds point opposite ways.

Polar, 1.60 D

Common mistakes

Polar bonds do not always make a polar molecule. CO₂, BF₃, CCl₄ and XeF₄ have polar bonds but no net dipole.
Judging the shape from the formula alone. SF₄ and XeF₄ look alike on paper, but the lone pairs give a seesaw and a square plane.
Using a flat drawing for a 3D molecule. CH₂Cl₂ drawn as a flat cross looks symmetric, but in 3D it is tetrahedral and polar.
Pointing the arrow the wrong way. In chemistry the bond dipole arrow points to the δ− (more electronegative) atom. Physics often uses the opposite sign.
Treating the cut-offs as exact. 0.4 and 1.7 are guides. Bond character changes smoothly with ΔEN.

The vector sum is a model in ΔEN units, used to decide whether the bond dipoles cancel. Real dipole moments (in debye) also depend on bond lengths and lone pairs; measured values are shown where known.

Show the dipoles in your own drawing. Draw structures, add δ+ and δ− labels and arrows in MolDraw, a free online chemical structure editor, and export a PNG or SVG.

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Frequently asked questions

How do you know if a bond is polar or nonpolar?

Subtract the Pauling electronegativities. A difference below 0.4 is nonpolar covalent, 0.4 to 1.7 is polar covalent and above 1.7 is usually called ionic. H–Cl: 3.16 − 2.20 = 0.96, so the bond is polar covalent with Cl δ−.

How do you calculate percent ionic character?

Pauling’s formula is % ionic = (1 − e^(−ΔEN²/4)) × 100. For H–Cl, ΔEN = 0.96, so % ionic = (1 − e^(−0.23)) × 100 = 20.6%. For H–F, ΔEN = 1.78 gives about 55%.

How do you know if a molecule is polar?

Find the shape with VSEPR, draw a dipole arrow for each polar bond and add them as vectors. If they cancel (symmetric shape with identical outer atoms and no lone pairs on the centre), the molecule is nonpolar. If not, it is polar. Water is bent, so its O–H dipoles do not cancel.

Why is CO2 nonpolar when its C=O bonds are polar?

CO2 is linear. The two C=O bond dipoles are equal and point in opposite directions, so they cancel and the net dipole moment is 0.

Why is CCl4 nonpolar but CHCl3 polar?

In tetrahedral CCl4 the four identical C–Cl dipoles cancel. In CHCl3 one Cl is replaced by H, so the dipoles no longer cancel. CHCl3 has a dipole moment of about 1.04 D.

Why is XeF4 nonpolar but SF4 polar?

XeF4 is square planar: its two lone pairs sit opposite each other and the four Xe–F dipoles cancel. SF4 is seesaw shaped because of one lone pair, so its S–F dipoles do not cancel.

Is cis- or trans-1,2-dichloroethene polar?

The cis isomer is polar (about 1.9 D) because both C–Cl dipoles point to the same side. The trans isomer is nonpolar (0 D) because the two C–Cl dipoles point in opposite directions and cancel.

Related chemistry tools

Read the full guide: Bond polarity and dipole moments explained.

Last updated: October 19, 2026