A molecule is polar if it has an uneven distribution of electrical charge, meaning one side is slightly negative and the other slightly positive. To figure this out, you need two pieces of information: whether the individual bonds in the molecule are polar, and whether the molecule’s shape allows those polar bonds to cancel each other out. A molecule with polar bonds can still be nonpolar overall if its geometry is symmetrical enough.
Step 1: Check the Bonds
Every bond between two different atoms has some degree of polarity. The question is whether it’s enough to matter. Polarity in a bond comes from one atom pulling on shared electrons harder than the other. This pulling power is measured on the Pauling electronegativity scale, and the bigger the difference between two bonded atoms, the more polar the bond.
Here are the key electronegativity values you’ll use most often:
- Hydrogen: 2.2
- Carbon: 2.55
- Nitrogen: 3.04
- Oxygen: 3.44
Subtract the smaller value from the larger to get the electronegativity difference. That number tells you what kind of bond you’re dealing with:
- 0.0 to 0.4: Nonpolar covalent (examples: C–H, C–C)
- 0.5 to 0.9: Slightly polar (examples: H–N, H–Cl)
- 1.0 to 1.3: Moderately polar (examples: C–O, S–O)
- 1.4 to 1.7: Highly polar (examples: H–O)
- 1.8 and above: Ionic territory (examples: H–F, NaF)
A practical shortcut: if a bond connects carbon or hydrogen to nitrogen, oxygen, fluorine, or chlorine, treat it as polar. Bonds between two identical atoms (like O–O or C–C) are always nonpolar. Bonds between carbon and hydrogen have a difference of only 0.35, so they’re treated as nonpolar for most purposes.
If a molecule contains no polar bonds at all, it’s nonpolar and you’re done. If it has exactly one polar bond, the molecule is polar and you’re also done. The more interesting cases are molecules with multiple polar bonds, because then geometry becomes the deciding factor.
Step 2: Determine the Molecule’s Shape
To find the shape, start by drawing the Lewis structure and counting the groups of electrons around the central atom. A “group” can be a single bond, a double bond, a triple bond, or a lone pair of electrons. Double and triple bonds each count as just one group.
The number of electron groups determines the basic geometry:
- 2 groups: Linear (atoms arranged in a straight line)
- 3 groups: Trigonal planar (flat triangle)
- 4 groups: Tetrahedral (3D pyramid shape)
- 5 groups: Trigonal bipyramidal
- 6 groups: Octahedral
But here’s the catch: lone pairs take up space around the central atom without being visible in the molecule’s shape. So the electron geometry and the molecular shape can be different. A molecule with four electron groups but one lone pair doesn’t look tetrahedral. It looks like a pyramid (trigonal pyramidal). You determine the shape by looking at only the positions of the atoms, ignoring where the lone pairs sit.
Step 3: See If the Polar Bonds Cancel Out
This is where many students get tripped up. A molecule can have very polar bonds and still be nonpolar overall, because the pulls from those bonds point in opposite directions and cancel out. Think of it like a tug-of-war: if equal forces pull from all sides, nothing moves.
Molecules with no lone pairs on the central atom and identical atoms attached to it are almost always nonpolar, regardless of how polar the individual bonds are. The symmetric charge distribution causes the bond polarities to cancel perfectly. These nonpolar-despite-polar-bonds shapes include:
- Linear (like CO₂, where two equally polar C=O bonds point in opposite directions)
- Trigonal planar (like BF₃, where three identical bonds fan out evenly at 120°)
- Tetrahedral (like methane, CH₄, with four identical bonds spaced evenly in 3D)
- Trigonal bipyramidal (like PCl₅)
- Octahedral (like SF₆)
- Square planar (like XeF₄, which has two lone pairs but they sit opposite each other, so the bond dipoles still cancel)
Molecules that are polar typically have lone pairs on the central atom, or they have different types of atoms bonded to the center, breaking the symmetry. Water is the classic example: it has two polar O–H bonds and two lone pairs, giving it a bent shape. The two bond polarities point in roughly the same direction instead of canceling, making water very polar. Ammonia (NH₃) is similar. It has three polar N–H bonds and one lone pair, creating a pyramidal shape where the polarities add up rather than cancel.
A useful mental trick: draw arrows along each polar bond, pointing from the less electronegative atom toward the more electronegative one. Imagine these arrows as physical forces pushing the molecule. If they’d push the molecule in some net direction, the molecule is polar. If they balance out perfectly, it’s nonpolar.
Quick-Reference Examples
Seeing a few examples side by side makes the pattern click:
- CO₂: Two polar C=O bonds, but the linear shape means they point in exactly opposite directions. Nonpolar.
- H₂O: Two polar O–H bonds in a bent shape. The pulls don’t cancel. Polar.
- CH₄: Four C–H bonds with a tiny electronegativity difference (0.35), arranged in a perfect tetrahedron. Nonpolar.
- CHCl₃ (chloroform): Three polar C–Cl bonds and one C–H bond. The different atoms break the tetrahedral symmetry. Polar.
- CCl₄: Four identical C–Cl bonds in a perfect tetrahedron. The polar bonds cancel. Nonpolar.
- NH₃: Three polar N–H bonds plus a lone pair making it pyramidal. Polar.
Notice how CCl₄ and CHCl₃ have nearly the same structure, but swapping even one atom from chlorine to hydrogen destroys the symmetry and flips the result from nonpolar to polar.
Testing Polarity in Practice
If you’re working in a lab rather than on paper, the simplest test is “like dissolves like.” Polar substances dissolve in polar solvents, and nonpolar substances dissolve in nonpolar solvents. Water (highly polar, with a polarity index of 10.2) won’t mix with hexane (nearly nonpolar, polarity index 0.1). If you drop an unknown substance into water and it dissolves readily, it’s likely polar. If it refuses to mix with water but dissolves easily in something like hexane or mineral oil, it’s nonpolar.
A classic demonstration of this: solid iodine barely dissolves in water, producing only a faint color change. Drop the same iodine into a nonpolar solvent like cyclohexane and it dissolves immediately, turning an intense purple. The iodine molecule (I₂) is nonpolar, two identical atoms sharing electrons equally, so it strongly prefers nonpolar company.
You can also observe this when two liquids refuse to mix. Water and carbon tetrachloride (CCl₄) form two distinct layers because one is polar and the other isn’t. No amount of shaking will make them combine into a single solution.
Molecules That Are Both Polar and Nonpolar
Some molecules are large enough to have a polar region and a nonpolar region in the same structure. Soap is the most familiar example. A soap molecule has a long hydrocarbon chain (a tail made of carbon and hydrogen, which is nonpolar) attached to a charged or highly polar “head group” (typically a carboxylate group with oxygen atoms). This dual nature is what makes soap useful: the nonpolar tail dissolves into grease and oil, while the polar head dissolves into water, pulling the grease away.
These dual-natured molecules are called amphiphilic. Phospholipids in your cell membranes work the same way. If you’re looking at a large organic molecule and trying to assess its polarity, look for oxygen, nitrogen, or charged groups (the polar parts) and long carbon-hydrogen chains (the nonpolar parts). The overall behavior of the molecule depends on which region dominates or on how you’re interacting with it.
The Dipole Moment as a Measure
Chemists quantify a molecule’s polarity using its dipole moment, measured in units called Debyes. Most polar molecules have dipole moments around 1 Debye, with highly polar molecules going higher. A dipole moment of exactly zero means the molecule is nonpolar. There’s no strict cutoff between “polar” and “nonpolar” on this scale. Instead, it’s a spectrum. Molecules with very small electronegativity differences across their bonds will have dipole moments close to zero and behave as nonpolar in practice, even if they’re technically not at zero.
You won’t typically calculate dipole moments by hand in a general chemistry course, but understanding that polarity is a spectrum rather than a binary switch helps explain why some molecules behave as “slightly polar” in certain situations. Acetone, for instance, sits in the middle of the polarity scale (polarity index 5.1) and can dissolve both moderately polar and moderately nonpolar substances, making it a useful general-purpose solvent.