Predicting polarity comes down to two things: whether the bonds in a molecule are polar, and whether the molecule’s shape allows those polar bonds to cancel out. If the bond polarities don’t cancel, the molecule is polar. If they do, it’s nonpolar. This straightforward process works for the vast majority of molecules you’ll encounter in general or organic chemistry.
Step 1: Find the Electronegativity Difference
Every atom has a measurable tendency to pull electrons toward itself in a bond. This is its electronegativity, and the most widely used scale is the Pauling scale. The bigger the difference between two bonded atoms, the more polar the bond. Here are the values you’ll use most often:
- Fluorine: 3.98 (the highest of any element)
- Oxygen: 3.44
- Nitrogen: 3.04
- Chlorine: 3.16
- Bromine: 2.96
- Carbon: 2.55
- Hydrogen: 2.20
- Sulfur: 2.58
- Sodium: 0.93
Subtract the smaller value from the larger one. That difference tells you the bond type:
- Below about 0.4: The bond is essentially nonpolar covalent. Carbon and hydrogen, for example, differ by only 0.35, which is why C–H bonds are treated as nonpolar in practice.
- Between roughly 0.4 and 1.5: The bond is polar covalent. Electrons are shared unevenly, creating partial positive and partial negative ends. An O–H bond (difference of 1.24) falls squarely here.
- Above about 1.5: The bond is ionic. Electrons are essentially transferred rather than shared. Sodium chloride has a difference of 2.23, which is a textbook ionic bond.
These thresholds aren’t sharp cutoffs. They’re guidelines that reflect a smooth spectrum from pure covalent to fully ionic. But for predicting molecular polarity, they give you a reliable starting point: if none of the bonds in a molecule are polar, the molecule is nonpolar, and you can stop here.
Step 2: Determine the Molecular Shape
A molecule can have polar bonds and still be nonpolar overall. That happens when the shape of the molecule arranges those polar bonds so they pull in opposite directions and cancel out. To figure out the shape, you need to count the groups of electrons around the central atom (bonding pairs and lone pairs) and apply VSEPR theory, which predicts that electron groups spread out as far apart as possible.
The shapes that matter most for polarity prediction:
- Linear (2 bonds, no lone pairs): Bond angles of 180°. If both bonds are identical, the dipoles cancel. CO₂ is the classic example.
- Trigonal planar (3 bonds, no lone pairs): Bond angles of 120°. If all three bonds are identical, the dipoles cancel. BF₃ is nonpolar for this reason.
- Tetrahedral (4 bonds, no lone pairs): Bond angles of about 109.5°. If all four bonds are identical, the dipoles cancel. CH₄ and CCl₄ are both nonpolar.
- Bent (2 bonds + 1 or 2 lone pairs): The lone pairs push the bonding pairs closer together, creating an asymmetric shape. Water is bent, and its dipoles do not cancel.
- Trigonal pyramidal (3 bonds + 1 lone pair): The lone pair sits on top, pushing the three bonds downward into a pyramid. Ammonia (NH₃) has this shape and is polar.
The key principle: lone pairs on a central atom almost always break the symmetry needed for dipoles to cancel. If your central atom has one or more lone pairs, the molecule is very likely polar.
Step 3: Add the Bond Dipoles as Vectors
Each polar bond creates a small dipole, a tiny arrow pointing from the less electronegative atom toward the more electronegative one. The molecule’s overall polarity is the vector sum of all those arrows. If they add up to zero, the molecule is nonpolar. If they don’t, it’s polar.
You don’t usually need to do formal vector math. For most molecules, you can reason it out visually. In CO₂, two equally strong C=O dipoles point in exactly opposite directions (180° apart), so they cancel perfectly. Net dipole: zero. In water, the two O–H dipoles point toward oxygen, but because the molecule is bent at about 104.5°, they reinforce each other instead of canceling. The net dipole moment of water is 1.84 D (debye), calculated as 2 × 1.5 D × cos(104.5°/2).
For tetrahedral molecules, the logic is the same. CCl₄ has four identical C–Cl dipoles arranged symmetrically, so they cancel completely. But CHCl₃ (chloroform) replaces one chlorine with hydrogen, breaking the symmetry. The bond dipoles no longer cancel, and the molecule is polar. This is a pattern worth remembering: identical atoms arranged symmetrically around a center cancel out, but mixing different atoms or adding lone pairs breaks the cancellation.
Common Molecules and Why They’re Polar or Not
Working through familiar examples helps build intuition faster than memorizing rules.
Nonpolar examples: CO₂ is linear and symmetric, so its two strong C=O dipoles cancel. Methane (CH₄) is tetrahedral with four nearly identical bonds. CCl₄ and CF₄ are also tetrahedral with identical bonds on every side. BF₃ is trigonal planar and symmetric. Propane and other hydrocarbons are nonpolar because C–H bonds have such a small electronegativity difference (0.35) that there’s essentially no dipole to worry about. Diatomic molecules made of the same element, like O₂, are nonpolar because two identical atoms pull on the electrons equally.
Polar examples: HF is a diatomic molecule with a large electronegativity difference (1.78), so it’s straightforwardly polar. Water is bent with two lone pairs on oxygen, giving it a strong net dipole. Ammonia (NH₃) is trigonal pyramidal with one lone pair, making it polar. Hydrogen cyanide (HCN) is linear, but nitrogen and hydrogen have different electronegativities, so the dipoles don’t cancel. Chloromethane (CH₃Cl) is tetrahedral, but carbon is bonded to both hydrogen and chlorine, creating an asymmetric arrangement where all the bond moments point toward the chlorine end. Methanol is polar because the –OH group creates an asymmetric charge distribution that the rest of the molecule can’t balance.
A Quick Decision Flowchart
When you need to predict polarity quickly, run through these questions in order:
- Are there any polar bonds? Check the electronegativity difference. If all bonds have a difference below about 0.4, the molecule is nonpolar. This covers all pure hydrocarbons.
- Does the central atom have lone pairs? If yes, the molecule is almost certainly polar. Lone pairs create asymmetry that prevents dipole cancellation.
- Are all the atoms around the central atom identical? If yes and there are no lone pairs, the molecule is nonpolar regardless of how polar the individual bonds are. This is why CCl₄ and BF₃ are nonpolar.
- Are the surrounding atoms different? If the central atom is bonded to a mix of different atoms (like CHCl₃), the molecule is polar because the different bond strengths can’t cancel symmetrically.
Why Polarity Matters: Predicting Solubility
The most practical payoff of predicting polarity is knowing what dissolves in what. The rule is simple: like dissolves like. Polar substances dissolve in polar solvents, and nonpolar substances dissolve in nonpolar solvents. Mixing polar and nonpolar usually fails.
Table salt (NaCl) dissolves readily in water because both are polar (ionic compounds count as extremely polar). Oil doesn’t dissolve in water because oil molecules are nonpolar hydrocarbons, while water is highly polar. Wax dissolves in hexane because both are nonpolar. This is also why grease comes off your hands with soap but not with water alone: soap molecules have both a polar end and a nonpolar end, bridging the gap between the two worlds.
The underlying reason is that molecules with similar charge distributions experience similar intermolecular forces. Polar molecules interact through dipole-dipole attractions and, when O–H or N–H bonds are present, hydrogen bonding. Nonpolar molecules rely on weaker London dispersion forces. When a polar solute is dropped into a nonpolar solvent, it can’t form the strong interactions it needs to stay dissolved, so it separates out. Liquids that dissolve in each other in all proportions are called miscible (like water and ethanol), while those that refuse to mix are immiscible (like water and oil).