Molecular polarity
A molecule is polar when its charge distribution has a nonzero electric dipole moment. Two ingredients are required: individual bonds must carry dipoles, and the three-dimensional geometry must leave a nonzero vector sum.
Bond polarity starts with unequal electron attraction
Electronegativity differences shift shared electron density. The more electronegative end is labeled δ− and the other δ+. These are partial charges, not formal oxidation states or full ionic charges.
The bond dipole moment depends on both separated charge and distance, so Δχ alone does not determine the measured dipole.
Geometry decides whether dipoles cancel
CO₂ has two strongly polar C=O bonds, but its linear symmetry makes the dipoles cancel. H₂O is bent, so the two O–H dipoles add to a net molecular dipole.
BF₃, CH₄, PF₅ and SF₆ provide other examples where high symmetry cancels equivalent bond dipoles.
Lone pairs often change both shape and polarity
Lone pairs affect the electron-domain geometry and can destroy the symmetry needed for cancellation. NH₃ is trigonal pyramidal rather than trigonal planar; SO₂ is bent rather than linear.
But a lone pair does not automatically guarantee a polar molecule. XeF₄ has two lone pairs yet is square planar with symmetric bond-dipole cancellation.
Polarity influences intermolecular behavior
Polar molecules interact through orientation-dependent dipole–dipole forces. If H is bonded to N, O or F, particularly strong directional hydrogen bonding may occur.
Polarity influences solubility, boiling point, dielectric behavior and how molecules orient in electric fields, but molecular size and dispersion forces remain important. “Polar = high boiling point” is not a universal rule.
Molecule selector
Measured dipole moments and limits
Dipole moments are measured quantities, commonly reported in debye (D). A zero dipole moment does not mean “no charge separation anywhere”; it can mean that local dipoles cancel exactly by symmetry.
Highly flexible molecules sample many conformations, so their effective polarity can depend on conformation and environment. Solvents can also reshape charge distribution.
Dipole moment as a measurable vector
For two separated charges ±q, the dipole-moment magnitude is μ = qr. Molecular charge distributions are continuous rather than two point charges, but the same vector idea survives: the dipole moment records how positive and negative charge are displaced in space.
Dipole moments are commonly reported in debye. HCl is polar because its electron distribution is shifted toward Cl. CO₂ has local C=O bond dipoles but a molecular dipole of zero by symmetry.
A molecule with μ = 0 can still have polarizable electron density and strong intermolecular dispersion forces. “Nonpolar” does not mean “electrostatically featureless.”
Symmetry gives a fast polarity test
| Geometry / example | Equivalent outer atoms? | Net dipole? |
|---|---|---|
| Linear CO₂ | Yes | No |
| Trigonal planar BF₃ | Yes | No |
| Tetrahedral CCl₄ | Yes | No |
| Bent H₂O | Yes, but geometry is not cancellation-symmetric | Yes |
| Trigonal pyramidal NH₃ | Yes, but lone pair breaks inversion/cancellation | Yes |
| Tetrahedral CH₂Cl₂ | No | Yes |
This is a vector-symmetry argument, not a memorized list. Ask whether the bond-dipole vectors are related by symmetry operations that force their sum to zero.
Polarity in liquids and solubility
Polar molecules can be stabilized by polar solvents through dipole interactions, and ions are especially stabilized by ion–dipole interactions. This motivates the familiar “like dissolves like” rule, but the rule is only qualitative.
Large nonpolar regions, hydrogen-bond donors/acceptors, molecular size and entropy all matter. Ethanol mixes with water because its OH group hydrogen-bonds strongly, while longer-chain alcohols become progressively less water-soluble as the hydrocarbon region grows.
Polarity therefore connects molecular geometry to bulk behavior, but predicting solubility requires the whole balance of intermolecular interactions.
Electrostatic-potential maps add spatial detail
Color-coded electrostatic-potential surfaces are often used to visualize where a molecule is relatively electron-rich or electron-poor. They can reveal local polar regions even in a molecule whose total dipole is zero.
For example, CO₂ has electron-rich oxygen ends and an electron-poor carbon center, yet the overall vector dipole cancels. Local electrostatic interactions can still matter when another molecule approaches one end.
This is a useful correction to the idea that “nonpolar” means uniform electron density. Molecular polarity is one global vector property of a much richer charge distribution.
Conformation can change a molecular dipole
Flexible molecules do not have one rigid geometry. Rotation about single bonds changes the directions of individual bond dipoles, so different conformations can have different net dipole moments.
In solution, temperature and solvent interactions alter the conformational population. The measured average dipole can therefore reflect an ensemble rather than one textbook drawing.
Exercises
BF₃
Each B–F bond is polar. Why is BF₃ nonpolar?
Solution
Its trigonal-planar symmetry makes the three equivalent bond-dipole vectors cancel.
NH₃ versus BF₃
Which is polar and what geometric difference matters?
Solution
NH₃ is polar because its trigonal-pyramidal shape lacks the planar threefold cancellation of BF₃.
XeF₄
Does the presence of lone pairs make XeF₄ polar?
Solution
No. The two lone pairs are opposite and the four Xe–F bonds form a symmetric square plane, so the bond dipoles cancel.