Home / Learn / Chemistry / Dipole polarityTool #L54 (formerly #1109)

Dipole polarity

Pick a molecule, see each bond dipole as an arrow from the electronegativity difference, and watch them add up or cancel on the 3D VSEPR shape. Swap outer atoms (CH₄ to CH₃Cl to CCl₄), compare with measured dipole moments, see why NF₃ is far less polar than NH₃, and test yourself.

Keep it one tap away for everyday use.

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Add up the bond dipoles

Level

Drag the picture to turn it.

Grey arrows: bond dipoles, pointing to the atom that pulls the electrons harder (the + end is at the tail). Red arrow: their sum, the dipole of the whole molecule.

Try this

  • CO₂ and H₂O both have two polar bonds. Turn each one: why is only water polar?
  • Start from CH₄ and change one H to Cl, then two, then three, then four. When does the red arrow vanish again?
  • Compare NH₃ and NF₃ in "With numbers". The bond dipoles are about the same size, but the measured values are very different. Read the lone pair note.
  • XeF₄ has two lone pairs but no dipole. Where are they?
  • In SF₆, change one F to Cl: the perfect octahedron is no longer balanced.

Check yourself

Is each molecule polar or non-polar? Think about the shape first. New questions every time; nothing is saved.

What it does

A bond between two different atoms is polar: the more electronegative atom pulls the shared electrons towards itself and gets a small negative charge (δ−), the other gets δ+. That bond has a dipole, drawn as an arrow with a plus sign at its tail, pointing to the δ− end. A whole molecule is polar only if its bond dipoles do not cancel when you add them as arrows (vectors), and that depends on the shape. This page draws every bond dipole on a 3D model you can turn, sizes it by the electronegativity difference (ΔEN), and draws their sum as one red arrow: the dipole of the molecule. No red arrow means non-polar.

It follows the Class 11 Chemistry chapter on chemical bonding and molecular structure (polar covalent bonds, dipole moment, why CO₂ and BF₃ are non-polar while H₂O and NH₃ are polar). Find the shape first with the VSEPR shape viewer tool, count the bonds and lone pairs with lewis structures, and see the symmetry that makes dipoles cancel in the point group.

How to use it

  1. Choose a Molecule. The verdict, polar or non-polar, appears with a one-line reason.
  2. Drag the picture to turn it. The grey arrows are the bond dipoles; the red one is their sum.
  3. Swap any outer H or halogen with the drop-downs under the menu. Turning CH₄ into CH₃Cl, then CH₂Cl₂, CHCl₃ and CCl₄ shows how the balance breaks and comes back.
  4. Switch Level to With numbers to see each ΔEN, the vector sum written out, the model’s net value and the measured dipole moment in debye (D).
  5. In Check yourself, decide whether four molecules are polar, then press Check.

Worked example: CO₂ against H₂O

In CO₂ each C=O bond has ΔEN = 3.44 − 2.55 = 0.89, pointing to O. The molecule is linear: the two arrows point in exactly opposite directions and cancel. Net dipole 0, so CO₂ is non-polar, even with two polar bonds.

In H₂O each O–H bond has ΔEN = 3.44 − 2.20 = 1.24, pointing to O. The two lone pairs on oxygen make the molecule bent, so both arrows lean the same way. Their sum points from between the H atoms towards O: the molecule is polar (measured 1.85 D).

Why NF₃ is less polar than NH₃

Both are pyramidal with one lone pair on N, and their bond dipoles are similar in size (ΔEN 0.84 and 0.94). But in NH₃ the N–H dipoles point towards N, the same way as the lone pair, so they add: 1.47 D. In NF₃ the N–F dipoles point away from N, against the lone pair, so they partly cancel: only 0.23 D. The simple arrow model leaves out the lone pair, so the tool adds a note whenever one helps or hinders.

Tips

  • Same outer atoms + a symmetric shape (linear, trigonal planar, tetrahedral, trigonal bipyramidal, square planar, octahedral) = non-polar.
  • Lone pairs on the central atom usually make a shape lopsided: bent, pyramidal, see-saw and T-shaped molecules are polar.
  • Different outer atoms (CH₃Cl, CHCl₃) break the balance even on a perfect tetrahedron.

FAQ

Are the numbers real dipole moments? The model value is only a relative size from ΔEN; it ignores lone pairs, bond lengths and real angles. The measured values in debye are real data-book values.

Is this an NCERT tool? No. It is a free fibub page made for this chapter; it runs on your device.

Where the numbers come from

  • Pauling electronegativities (H 2.20, C 2.55, N 3.04, O 3.44, F 3.98, Cl 3.16 and so on) are the standard published values. Bond dipoles as ΔEN along the bond, added as vectors over the ideal VSEPR shape, are a relative teaching model written fresh for fibub, not debye values.
  • Measured dipole moments (H₂O 1.85 D, NH₃ 1.47 D, NF₃ 0.23 D, HCl 1.11 D, CH₃Cl 1.87 D, CHCl₃ 1.04 D, SO₂ 1.63 D and others) are standard data-book gas-phase values. Polar and non-polar verdicts checked against the textbook outcomes for all 19 molecules.

Standard scientific values, such as constants, atomic weights and melting points, are facts found in every textbook; we name where ours were checked. If you are an author or publisher and want a value, a credit or a table changed or removed, tell us and we will act quickly.

Last updated 2026-11-09.

Change log for Dipole polarity

  1. Add #1109 Dipole moment and polarity Details on GitHub

Every change to fibub

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