By Jess Kroll, Updated Mar 24, 2022
When atoms form covalent bonds, they share electrons to create a stable molecule. If one atom has a stronger pull on the shared electrons, it gains a partial negative charge while the other atom carries a partial positive charge. This unequal charge distribution defines a polar molecule, whereas an even distribution results in a non‑polar molecule. To assess whether a specific atom is polar or non‑polar, follow these steps.
First, determine whether the bond is covalent or ionic. Ionic bonds form between ions—atoms that have gained or lost electrons and carry a net charge. Since ionic bonds involve charged species, the atoms themselves are not described as polar or non‑polar in the context of molecular polarity. In contrast, covalent bonds involve neutral atoms sharing electrons, and only atoms in such bonds can exhibit polarity.
Next, look at the elements that constitute the molecule. Homonuclear bonds (e.g., N₂, O₂, O₃) share electrons equally, making the atoms non‑polar. Heteronuclear bonds—those between different elements—tend to be polar because the atoms have differing electronegativities. Examples include CO₂, where the carbon‑oxygen bonds are polar, and H₂O, where the hydrogen‑oxygen bonds are strongly polar. If a molecule contains more than one element, the atoms involved in heteronuclear bonds are typically polar.
Finally, evaluate the overall shape of the molecule. Symmetric arrangements can cancel out individual dipole moments, rendering the entire molecule non‑polar even if its constituent atoms are polar. For instance, CO₂ is linear and non‑polar because its two polar bonds are mirror images. Asymmetric molecules, such as water (H₂O), possess a net dipole moment and are classified as polar. This step is crucial for predicting the physical properties of the substance.
These principles are foundational in chemistry education and are used by scientists to predict reactivity, solubility, and other key characteristics of compounds.
