Molecular polarity describes a separation of electric charge, resulting in an electric dipole moment with positive and negative ends. Lone pairs are non-bonding electron pairs that reside on an atom and are not shared in a covalent bond. The relationship between lone pairs and polarity is complex because polarity depends heavily on the molecule’s overall three-dimensional shape, which is often influenced by these unshared electrons.
Defining Molecular Polarity
A molecule is polar when it has an uneven distribution of electron density. This unevenness starts at the bond level, where a difference in electronegativity creates a bond dipole. The more electronegative atom pulls the shared electrons closer, acquiring a partial negative charge, while the other atom acquires a partial positive charge.
The polarity of the entire molecule is determined by the net molecular dipole moment, which is the vector sum of all individual bond dipoles. If bond dipoles are arranged symmetrically, they cancel out, resulting in a nonpolar molecule. For instance, linear carbon dioxide (\(\text{CO}_2\)) and tetrahedral carbon tetrachloride (\(\text{CCl}_4\)) are nonpolar because their symmetrical geometries cause the bond dipoles to cancel completely. Molecular polarity requires two conditions: the presence of at least one polar bond and an asymmetrical geometry that prevents cancellation.
How Lone Pairs Influence Molecular Geometry
Lone pairs are non-bonding valence electrons that significantly influence a molecule’s three-dimensional arrangement. The Valence Shell Electron Pair Repulsion (VSEPR) theory explains this influence, stating that all electron groups—both bonding and lone pairs—repel each other to maximize distance.
Lone pairs occupy a larger volume of space than bonding pairs because they are held only by the central atom’s nucleus. This results in lone pairs exerting a stronger repulsive force, which physically pushes bonding pairs closer together. This distortion reduces the molecule’s symmetry and decreases the bond angles compared to an ideal geometry.
For example, methane (\(\text{CH}_4\)) has a symmetrical tetrahedral angle of \(109.5^\circ\). When the central atom has lone pairs, this symmetry is broken, leading to asymmetrical shapes like trigonal pyramidal or bent.
Lone Pairs as a Cause of Polarity
Lone pairs often directly cause a molecule’s polarity because the resulting asymmetrical shape prevents individual bond dipoles from canceling. This lack of vector cancellation creates a net molecular dipole moment.
Water (\(\text{H}_2\text{O}\)) is a classic example, possessing two bonding pairs and two lone pairs on the central oxygen atom. The lone pairs push the two hydrogen atoms down, forcing the molecule into a bent geometry. The polarity of the oxygen-hydrogen bonds, combined with the bent shape, results in a large net dipole moment pointing toward the oxygen atom.
Similarly, ammonia (\(\text{NH}_3\)) has three bonding pairs and one lone pair on the central nitrogen atom. The lone pair repels the bonding pairs, creating a trigonal pyramidal shape. The bond dipoles point toward the more electronegative nitrogen atom, and their vectors combine to create a strong net dipole moment.
Polar Molecules That Do Not Contain Central Lone Pairs
While lone pairs on a central atom frequently cause a molecule to be polar, they are not a prerequisite for polarity. The true determinant of molecular polarity is asymmetry, which can arise in molecules that lack a central atom entirely.
Simple diatomic molecules illustrate this principle. For molecules like hydrogen chloride (\(\text{HCl}\)) or hydrogen fluoride (\(\text{HF}\)), polarity exists purely due to the difference in electronegativity between the two bonded atoms. Since there is no central atom, the electron density is pulled toward the more electronegative atom, creating a net dipole moment along the bond axis.
Polarity can also exist when the central atom has no lone pairs, but the bonded atoms are not identical. For example, chloromethane (\(\text{CH}_3\text{Cl}\)) has a central carbon atom with no lone pairs and a tetrahedral geometry. However, because the four bonded atoms are three hydrogens and one chlorine, the bond dipoles do not cancel. The strong pull of the chlorine atom creates a net dipole moment, making the molecule polar despite the lack of central lone pairs.