Molecular polarity describes how the electric charge is distributed across a molecule, which is determined by the sharing of electrons between its atoms. When electrons are shared unequally, one part of the molecule develops a slight negative charge while another part develops a slight positive charge, leading to a polar molecule. Understanding a molecule’s polarity is important because this characteristic dictates many of its physical properties, such as its solubility, melting point, and boiling point. For instance, highly polar molecules tend to dissolve well in other polar liquids, following the “like dissolves like” principle, which explains why water and oil do not mix.
Assessing Polarity at the Bond Level
The first step in determining molecular polarity involves examining the individual bonds within the structure. Bond polarity is governed by a concept called electronegativity, which is the measure of an atom’s ability to attract a shared pair of electrons toward itself. Atoms like fluorine and oxygen have high electronegativity values, meaning they exert a strong pull on bonding electrons. By calculating the difference in electronegativity (\(\Delta EN\)) between two bonded atoms, the resulting bond can be classified as nonpolar covalent, polar covalent, or ionic.
When the electronegativity difference is very small, typically less than 0.4 on the Pauling scale, the electrons are shared nearly equally, creating a nonpolar covalent bond. A moderate difference in electronegativity, usually between 0.4 and 1.8, results in a polar covalent bond, where electrons spend more time near the more electronegative atom. This unequal sharing establishes a separation of charge, known as a bond dipole moment. The more attractive atom gains a partial negative charge (\(\delta-\)) and the other atom gains a partial positive charge (\(\delta+\)). A very large difference, exceeding approximately 1.8, means the electron is essentially transferred, resulting in an ionic bond.
The Influence of Molecular Geometry
Having established the polarity of the individual bonds, the next consideration is the molecule’s three-dimensional shape. Even if a molecule contains multiple polar bonds, its overall polarity can still be nonpolar if the molecule possesses a high degree of symmetry. This occurs because the individual bond dipole moments are vector quantities and can effectively cancel each other out. The physical arrangement of the atoms dictates whether these vector forces will oppose one another perfectly.
In a linear molecule like carbon dioxide (\(\text{CO}_2\)), for example, the oxygen atoms are more electronegative than the central carbon, creating two equal bond dipoles pointing outward. Because the molecule is straight, the two dipoles pull with equal strength in opposite directions, resulting in a net cancellation that makes the entire molecule nonpolar. Similarly, in a trigonal planar molecule like boron trifluoride (\(\text{BF}_3\)), the three polar \(\text{B-F}\) bond dipoles are arranged symmetrically. This symmetrical arrangement ensures the vector sum of the three equal dipoles is zero, classifying \(\text{BF}_3\) as a nonpolar molecule. The same principle applies to symmetrical shapes like tetrahedral (\(\text{CH}_4\)) and octahedral (\(\text{SF}_6\)), where identical surrounding atoms and the absence of lone pairs lead to the cancellation of all bond dipoles.
Final Determination of Molecular Polarity
The final determination of a molecule’s polarity synthesizes the bond-level polarity with the molecular geometry by considering the net dipole moment. The net dipole moment is the overall measure of charge separation across the entire molecule, calculated as the vector sum of all individual bond dipoles. If this vector sum is zero, the molecule is classified as nonpolar. A non-zero net dipole moment indicates a polar molecule, meaning there is an overall separation of charge.
Asymmetry is the main factor that prevents bond dipoles from canceling, leading to a polar molecule. This asymmetry can arise in two primary ways: either by bonding different atoms to the central atom or by the presence of non-bonding lone pairs of electrons. A molecule like chloromethane (\(\text{CH}_3\text{Cl}\)) is tetrahedral, but the presence of the \(\text{C-Cl}\) bond alongside the \(\text{C-H}\) bonds breaks the symmetry, resulting in a net dipole moment. The presence of lone pairs on the central atom is the other common cause of asymmetry, as they distort the molecular shape.
Water (\(\text{H}_2\text{O}\)) is a classic example, as its two lone pairs on the central oxygen atom force the \(\text{O-H}\) bonds into a bent shape. This bent geometry means the two polar \(\text{O-H}\) bond dipoles do not oppose each other and instead combine to create a significant net dipole moment. Similarly, in ammonia (\(\text{NH}_3\)), the single lone pair on the nitrogen atom distorts the shape into a trigonal pyramidal arrangement, which prevents the \(\text{N-H}\) bond dipoles from canceling out.