Dipole Moments: Why Some Molecules Are Polar

Hold a charged plastic comb next to a thin stream of water from a tap and the stream bends towards it. Try the same with a stream of hexane (in a safe lab setting) and almost nothing happens. The difference is that water molecules are polar: each has a positive end and a negative end, and they turn to line up with the comb’s electric field. The quantity that measures this separation of charge is the dipole moment. This article explains what it is, how to predict it from a molecule’s bonds and shape, and why it matters.

From polar bonds to dipoles

When two atoms with different electronegativities share electrons, the more electronegative atom pulls the pair towards itself, gaining a partial negative charge (δ−) and leaving the other atom with a partial positive charge (δ+). This is a polar bond (see bond polarity).

A pair of equal and opposite charges separated by a distance is called an electric dipole. Each polar bond is a small dipole — a bond dipole.

Defining the dipole moment

The dipole moment, μ, measures the size of a dipole:

μ = q × d

where q is the magnitude of the separated charge and d is the distance between the charges.

• Dipole moment is a vector: it has both size and direction. By the usual chemistry convention it’s drawn as an arrow pointing from the δ+ end to the δ− end, with a small cross at the positive end (physicists often draw it the other way — just be consistent).
• The SI unit is the coulomb metre (C m), but chemists usually use the debye (D), named after Peter Debye. 1 D ≈ 3.336 × 10⁻³⁰ C m.

For a sense of scale: one full electron charge (1.602 × 10⁻¹⁹ C) separated by 100 pm (10⁻¹⁰ m) gives a dipole of 1.602 × 10⁻²⁹ C m, which is 4.80 D.

Worked example: how ionic is HCl?

The measured dipole moment of HCl is 1.08 D, and its bond length is 127 pm. If HCl were fully ionic (H⁺ and Cl⁻), its dipole moment would be:

μ(ionic) = 1.602 × 10⁻¹⁹ C × 1.27 × 10⁻¹⁰ m = 2.03 × 10⁻²⁹ C m = 2.03 × 10⁻²⁹ ÷ 3.336 × 10⁻³⁰ ≈ 6.10 D

So the percentage ionic character ≈ 1.08 ÷ 6.10 × 100 ≈ 18 %. In other words, the partial charges in HCl are about 0.18 of an electron’s charge. That’s a direct, measured way to see that HCl is a polar covalent molecule, not an ionic one.

Molecular dipoles: adding the arrows

A molecule with several polar bonds has an overall dipole moment equal to the vector sum of its bond dipoles (plus contributions from lone pairs). Because vectors have direction, bond dipoles can reinforce each other or cancel out. That’s why shape is essential (see VSEPR and molecular geometry).

Cancelling dipoles: non-polar molecules

• CO₂ is linear: O=C=O. Each C=O bond is polar, but the two bond dipoles point in exactly opposite directions and cancel. μ = 0.
• BF₃ is trigonal planar: three B–F dipoles at 120° to each other cancel.
• CH₄ and CCl₄ are tetrahedral: four identical bond dipoles cancel.
• SF₆ is octahedral: all six S–F dipoles cancel.

Symmetrical molecules with identical surrounding atoms and no lone pairs on the central atom are always non-polar.

Reinforcing dipoles: polar molecules

• H₂O is bent (104.5°). The two O–H bond dipoles both point partly towards oxygen, and the two lone pairs on oxygen add to the effect. μ = 1.85 D.
• NH₃ is trigonal pyramidal. The three N–H dipoles point partly towards nitrogen, reinforced by the lone pair. μ = 1.47 D.
• CHCl₃ (chloroform): replacing one Cl in CCl₄ with H breaks the symmetry. μ ≈ 1.0 D.

A tricky case: NF₃

Nitrogen trifluoride has the same shape as ammonia, and N–F bonds (Δχ = 0.94) are more polar than N–H bonds (Δχ = 0.84). You might expect NF₃ to be more polar than NH₃. In fact, NF₃’s dipole moment is only about 0.23 D. In NH₃, the bond dipoles point towards nitrogen, in the same direction as the lone pair’s contribution. In NF₃, the bond dipoles point away from nitrogen, towards the fluorines — opposite to the lone pair’s contribution — so they largely cancel. It’s a great example of why you have to think about direction, not just size.

Isomers with different polarities

1,2-Dichloroethene exists as two isomers. In the cis isomer, both Cl atoms are on the same side, and their C–Cl dipoles add up: μ ≈ 1.9 D. In the trans isomer, they’re on opposite sides and cancel: μ = 0. The two isomers have different boiling points (about 60 °C for cis and 48 °C for trans), largely because of this difference.

Table of dipole moments

Molecule	Shape	Dipole moment (D)
H₂, O₂, N₂	Linear (identical atoms)	0
CO₂	Linear	0
CH₄, CCl₄	Tetrahedral	0
HF	Linear	1.82
HCl	Linear	1.08
HBr	Linear	0.82
HI	Linear	0.44
H₂O	Bent	1.85
NH₃	Trigonal pyramidal	1.47
NF₃	Trigonal pyramidal	0.23
CHCl₃	Tetrahedral (distorted)	about 1.0
CH₃Cl	Tetrahedral (distorted)	1.87
Propanone (acetone)	Trigonal planar at C=O	2.88

The hydrogen halides show a clear trend: dipole moments fall from HF to HI, as the electronegativity difference falls.

How dipole moments are measured

Dipole moments are measured by placing a substance between two charged plates. Polar molecules tend to line up with the electric field, which increases the amount of charge the plates can store (the capacitance). Measuring this at different temperatures gives the dipole moment. Microwave spectroscopy of gases, which measures how molecules rotate in electric fields, gives very precise values too.

Why dipole moments matter

Boiling points and intermolecular forces

Polar molecules attract each other through permanent dipole–dipole forces: the δ+ end of one molecule attracts the δ− end of another. These add to the London forces present in all molecules, so polar molecules usually have higher boiling points than non-polar molecules of similar size (see intermolecular forces).

Compare propane (C₃H₈, M = 44, μ ≈ 0.08 D, boiling point −42 °C) with ethanal (CH₃CHO, M = 44, μ ≈ 2.7 D, boiling point 20 °C). Similar mass and size, very different boiling points.

Solubility

“Like dissolves like.” Polar solvents such as water dissolve polar and ionic substances, because the solvent’s dipoles can surround and stabilise charged or polar particles. Non-polar substances such as oils dissolve better in non-polar solvents (see polar vs non-polar molecules).

Microwave ovens

Microwaves make polar water molecules repeatedly rotate as they try to line up with the rapidly changing electric field. That motion is transferred to neighbouring molecules as heat. Non-polar substances, such as dry plastics and glass, don’t heat up as much in a microwave.

Common misconceptions

• “Any molecule with polar bonds is polar.” Symmetry can cancel the bond dipoles (CO₂, CCl₄).
• “More polar bonds always mean a bigger dipole moment.” Direction matters — NF₃ versus NH₃.
• “Dipole moment and charge are the same.” Dipole moment depends on both charge and distance.
• “Lone pairs don’t affect polarity.” They can add to or reduce the molecular dipole.

Key takeaways

• A dipole moment measures charge separation: μ = q × d, usually in debyes (1 D ≈ 3.336 × 10⁻³⁰ C m).
• It’s a vector; the molecular dipole is the vector sum of bond dipoles and lone-pair effects.
• Symmetrical shapes cancel dipoles (CO₂, BF₃, CCl₄); bent or pyramidal shapes usually don’t (H₂O, NH₃).
• Comparing a measured dipole with a fully ionic value gives the percent ionic character (HCl ≈ 18 %).
• Dipole moments affect boiling points, solubility and microwave heating.