How to Determine If a Molecule Is Polar: 3 Simple Steps

To determine if a molecule is polar, you need to check two things: the shape of the molecule and the difference in electronegativity between its atoms. A molecule is polar if it has a net dipole moment, meaning one end carries a partial positive charge and the other end carries a partial negative charge. You can work this out in three steps: draw the Lewis structure, determine the molecular geometry, and check if the bond dipoles cancel.

Step 1: Draw the Lewis Structure and Find Bond Dipoles

Start by drawing the Lewis structure of the molecule. This shows how the atoms are connected and where the valence electrons are placed. Once you have the structure, identify any bonds between atoms with different electronegativities. Electronegativity is a measure of how strongly an atom attracts electrons in a bond. The difference in electronegativity between two bonded atoms creates a bond dipole, which is a vector pointing toward the more electronegative atom.

For example, in a water molecule (H₂O), the oxygen atom is more electronegative than the hydrogen atoms. Each O–H bond has a bond dipole pointing toward oxygen. In carbon dioxide (CO₂), the oxygen atoms are more electronegative than carbon, so each C=O bond has a dipole pointing toward oxygen. At this stage, you have identified the individual bond dipoles, but you cannot yet say whether the whole molecule is polar. That depends on the shape.

A quick reference: if the electronegativity difference is less than about 0.4, the bond is usually considered nonpolar covalent. Between 0.4 and 1.7, it is polar covalent. Above 1.7, it is often ionic. You can find electronegativity values on a periodic table or from a reliable source like IUPAC.

Step 2: Determine the Molecular Geometry

The shape of the molecule determines whether the bond dipoles add up to a net dipole or cancel out. To find the shape, use the VSEPR (Valence Shell Electron Pair Repulsion) theory. This theory states that electron pairs around a central atom repel each other and arrange themselves as far apart as possible. The molecular geometry is the arrangement of the atoms only, ignoring lone pairs of electrons.

Count the number of bonding groups (atoms bonded to the central atom) and lone pairs on the central atom. This gives you the electron domain geometry, which then tells you the molecular geometry. Common geometries include:

• Linear (2 bonding groups, 0 lone pairs) – e.g., CO₂, BeCl₂
• Trigonal planar (3 bonding groups, 0 lone pairs) – e.g., BF₃
• Bent (2 bonding groups, 1 lone pair) – e.g., SO₂
• Tetrahedral (4 bonding groups, 0 lone pairs) – e.g., CH₄
• Trigonal pyramidal (3 bonding groups, 1 lone pair) – e.g., NH₃
• Bent (2 bonding groups, 2 lone pairs) – e.g., H₂O

You can find detailed explanations of VSEPR theory on Khan Academy. The key point is that symmetrical shapes often lead to nonpolar molecules, while asymmetrical shapes often lead to polar molecules.

Worked Example: Water (H₂O)

Water has a central oxygen atom with two O–H bonds and two lone pairs. The electron domain geometry is tetrahedral, but the molecular geometry is bent because the lone pairs are not counted as atoms. The bond dipoles in water point toward oxygen, and because the molecule is bent, these dipoles do not cancel. They add up to a net dipole pointing from the hydrogen side (partial positive) to the oxygen side (partial negative). Water is polar.

Step 3: Check If the Bond Dipoles Cancel

Now that you know the shape, you can decide whether the molecule has a net dipole moment. Draw the bond dipoles as vectors, and see if they point in opposite directions and are equal in magnitude. If they cancel, the molecule is nonpolar. If they do not cancel, the molecule is polar.

Symmetrical vs. Asymmetrical Molecules

In a symmetrical molecule, the bond dipoles often cancel. For example, carbon dioxide (CO₂) is linear: O=C=O. The two C=O bond dipoles point toward the oxygen atoms, which are in opposite directions. They are equal in magnitude, so they cancel exactly. CO₂ is nonpolar, even though each bond is polar.

In contrast, water is bent, so the two O–H dipoles do not point opposite each other. They combine to give a net dipole. Similarly, ammonia (NH₃) is trigonal pyramidal. The three N–H bond dipoles all point toward nitrogen, but because the molecule is not flat, the dipoles do not cancel. Ammonia is polar.

Methane (CH₄) is tetrahedral and symmetrical. The four C–H bond dipoles are equal in magnitude and point toward the corners of a tetrahedron. They cancel exactly, so methane is nonpolar. Note that the C–H bond is only weakly polar (electronegativity difference of about 0.4), but even if it were strongly polar, the symmetry would still cause cancellation.

Molecules with Lone Pairs

Lone pairs themselves do not create bond dipoles, but they affect the molecular geometry. A molecule with lone pairs on the central atom is often polar because the lone pairs push the bonding groups closer together, breaking symmetry. For example, sulfur dioxide (SO₂) has a bent shape (2 bonding groups, 1 lone pair on sulfur). The S=O bond dipoles do not cancel, so SO₂ is polar.

There are exceptions. Carbon dioxide has no lone pairs on carbon, but it is linear and nonpolar. Xenon tetrafluoride (XeF₄) has a square planar shape (4 bonding groups, 2 lone pairs on xenon). The bond dipoles cancel because the fluorine atoms are opposite each other, so XeF₄ is nonpolar. You need to check the specific geometry rather than assuming all molecules with lone pairs are polar.

Comparing Molecules with the Same Shape but Different Polarity

Two molecules can have the same molecular geometry but different polarity if their outer atoms are not identical. Consider a molecule like CH₂Cl₂ (dichloromethane). It has a tetrahedral shape, but the outer atoms are not all the same. Two hydrogen atoms and two chlorine atoms are bonded to carbon. The C–Cl bonds are polar (chlorine is more electronegative than carbon), while the C–H bonds are only weakly polar. Because the chlorine atoms are adjacent to each other, the two C–Cl dipoles do not cancel. They add up to a net dipole pointing from the hydrogen side to the chlorine side. CH₂Cl₂ is polar.

If the chlorine atoms were opposite each other, the dipoles would cancel. But in tetrahedral geometry with two identical polar groups and two identical nonpolar groups, the molecule is polar if the polar groups are adjacent and nonpolar if they are opposite. This is a useful rule for organic molecules.

Key Takeaways

• A molecule is polar if it has a net dipole moment, which depends on both bond polarity and molecular shape.
• Step 1: Draw the Lewis structure and identify bond dipoles from electronegativity differences.
• Step 2: Use VSEPR theory to determine the molecular geometry.
• Step 3: Check if the bond dipoles cancel. Symmetrical shapes (linear, tetrahedral, trigonal planar with identical outer atoms) are usually nonpolar. Asymmetrical shapes (bent, trigonal pyramidal) are usually polar.
• Lone pairs affect geometry but do not guarantee polarity; always check the vector sum of bond dipoles.
• For a reliable reference on electronegativity and bonding, consult resources like ChemGuide.