Molecular geometries that are asymmetric produce polar molecules, while highly symmetric geometries do not. Bent, trigonal pyramidal, seesaw, and T-shaped geometries are always polar when they contain polar bonds, because their shapes prevent bond dipoles from canceling out. Symmetric geometries like linear (with identical terminal atoms), trigonal planar, tetrahedral, trigonal bipyramidal, square planar, and octahedral are nonpolar when all surrounding atoms are the same, because their bond dipoles cancel perfectly.
The key principle is straightforward: polarity depends on both bond polarity and shape. A molecule needs at least one polar bond, and those bond dipoles must add up to a nonzero net dipole rather than canceling each other out.
Why Shape Determines Polarity
Every polar bond has a small dipole, a tiny arrow of charge pointing from the less electronegative atom toward the more electronegative one. When you place several of these arrows in three-dimensional space, they can either reinforce each other or cancel out, depending on the geometry. Carbon dioxide is the classic example of cancellation: each C–O bond is quite polar, but because the molecule is linear (180° apart), the two dipoles point in exactly opposite directions and sum to zero. Water has the same two polar bonds to oxygen, but its bent shape (about 104.5°) means the dipoles point in roughly the same direction. They add up instead of canceling, giving water a dipole moment of 1.85 Debye in the gas phase.
This is why you can’t determine polarity from bond polarity alone. You have to know the shape.
Geometries That Are Always Polar
Certain molecular shapes are inherently lopsided. As long as the bonds in these molecules are polar, the molecule will be polar, because there’s no way for the dipoles to cancel.
- Bent (V-shaped): Two bonding pairs and one or two lone pairs on the central atom. Water (H₂O) and sulfur dioxide (SO₂) are examples. The bond angle is less than 180°, so the dipoles always add up to a net dipole.
- Trigonal pyramidal: Three bonding pairs and one lone pair. Ammonia (NH₃) is the textbook case. The three N–H dipoles all point upward toward nitrogen, and because the lone pair sits on top, nothing opposes them. The vectors add to a dipole pointing straight along the vertical axis.
- Seesaw (sawhorse): Four bonding pairs and one lone pair, derived from trigonal bipyramidal geometry. Sulfur tetrafluoride (SF₄) is a common example. The lone pair creates asymmetry that prevents full cancellation.
- T-shaped: Three bonding pairs and two lone pairs. Chlorine trifluoride (ClF₃) has this shape. The uneven distribution of bonds and lone pairs guarantees a net dipole.
The common thread is lone pairs on the central atom. Lone pairs take up space and push bonding pairs into asymmetric arrangements. Textbooks describe them as “stereochemically active,” meaning they don’t form bonds but they dictate the shape of the molecule. That shape-distorting effect is exactly what makes these geometries polar.
Geometries That Are Nonpolar (With a Catch)
Highly symmetric geometries cancel out their bond dipoles when all the terminal atoms are identical. These include:
- Linear (two identical bonds): CO₂, BeCl₂. Dipoles point in opposite directions and cancel to zero.
- Trigonal planar: BF₃. Three identical bonds spaced 120° apart. The dipole vectors sum to zero.
- Tetrahedral: CH₄, CCl₄. Four identical bonds arranged symmetrically. All dipoles cancel.
- Trigonal bipyramidal: PF₅. Five identical bonds cancel perfectly.
- Square planar: XeF₄. Four bonds at 90° with two lone pairs directly opposite each other. The symmetry produces a net dipole of zero.
- Octahedral: SF₆. Six identical bonds cancel completely.
Here’s the catch: these geometries are only nonpolar when all the surrounding atoms are the same. Replace even one atom with something different, and the symmetry breaks. Chloromethane (CH₃Cl) has a tetrahedral geometry, just like methane. But because the C–Cl bond has a larger dipole than the C–H bonds, the dipoles no longer cancel perfectly. Chloromethane is polar. The same principle applies to any symmetric geometry: swap in a different atom or group, and polarity can emerge.
How Electronegativity Fits In
Before shape even matters, there has to be a polar bond in the molecule. A bond is considered polar covalent when the electronegativity difference between the two atoms is about 0.5 or greater on the Pauling scale. Below 0.4, the bond is essentially nonpolar covalent (think C–H or C–C). From 0.5 to 0.9, it’s slightly polar (H–Cl). From 1.0 to 1.3, moderately polar (C–O). At 1.4 to 1.7, highly polar (H–O). Beyond about 1.8, the bond starts to become ionic rather than covalent.
A molecule made entirely of nonpolar bonds, like methane with its C–H bonds at an electronegativity difference of only 0.4, is nonpolar regardless of shape. In practice, though, most bonds in chemistry have at least some polarity, so shape usually ends up being the deciding factor.
How to Determine Polarity Step by Step
If you’re working through a specific molecule, here’s the process. First, check whether the molecule contains any polar bonds using electronegativity values. If every bond in the molecule is nonpolar, the molecule is nonpolar, and you’re done. If there’s exactly one polar bond, the molecule is polar, no further analysis needed.
If there are two or more polar bonds, draw the Lewis structure and count the bonding pairs and lone pairs around the central atom. Use VSEPR theory to determine the molecular geometry. Then mentally (or on paper) add up the bond dipole vectors based on the three-dimensional arrangement. If they cancel to zero, the molecule is nonpolar. If they don’t, it’s polar.
In practice, this means you can often shortcut the process: if the geometry is bent, trigonal pyramidal, seesaw, or T-shaped, and the bonds are polar, the molecule is polar. If the geometry is one of the symmetric shapes and all surrounding atoms are identical, the molecule is nonpolar. The only cases that require careful vector addition are symmetric geometries with mixed terminal atoms, like CHCl₃ or CH₂Cl₂, where you need to assess whether the different bond dipoles still manage to cancel.
Quick Reference
- Always polar (with polar bonds): Bent, trigonal pyramidal, seesaw, T-shaped
- Nonpolar only when all terminal atoms match: Linear (2 bonds), trigonal planar, tetrahedral, trigonal bipyramidal, square planar, octahedral
- Always nonpolar: Any molecule with only nonpolar bonds, or any diatomic molecule of the same element (H₂, O₂, N₂)
The underlying rule never changes: polarity requires both a polar bond and an asymmetric arrangement. Geometry is what tips the balance.