Polarity describes the separation of electrical charge within a molecule or its bonds, resulting in a distinct positive end and a distinct negative end, known as a dipole. This charge asymmetry is fundamental to understanding how substances interact with one another, forming the basis for nearly all chemical and biological processes. A molecule that possesses this charge separation is termed a polar molecule, acting much like a tiny bar magnet. The presence of a dipole means the molecule has an electric dipole moment, a quantitative measure of the magnitude and direction of the charge separation within the structure.
The Origin of Polarity: Unequal Electron Sharing
The root of polarity lies in a property of atoms called electronegativity, which is a measure of an atom’s attraction for shared electrons in a chemical bond. When two atoms form a covalent bond, they share a pair of electrons. If the two bonded atoms have an identical electronegativity, the sharing is perfectly equal, creating a nonpolar covalent bond.
However, if the atoms have a measurable difference in their electronegativity, the shared electrons are pulled closer to the more attractive atom. This unequal tug-of-war causes the more electronegative atom to acquire a partial negative charge (\(delta-\)). The less electronegative atom is then left with a partial positive charge (\(delta+\)), resulting in a polar covalent bond. The greater the difference in electronegativity between the two atoms, the more polar the bond becomes.
How Molecular Shape Determines Polarity
The overall polarity of an entire molecule is not determined simply by the presence of polar bonds; it depends on the molecule’s three-dimensional shape, which dictates whether the individual bond dipoles cancel each other out. Each polar bond generates a dipole moment, which can be visualized as a vector pointing from the partial positive charge toward the partial negative charge. The total polarity of the molecule is the vector sum of all these individual bond dipoles, known as the net dipole moment.
A molecule with highly polar bonds can still be nonpolar if its structure is highly symmetrical, causing the dipole moments to pull equally in opposite directions and cancel out. For instance, carbon dioxide (\(text{CO}_2\)) has two carbon-oxygen polar bonds, but its linear shape means the two bond dipoles are equal in magnitude and point in exactly opposite directions, resulting in a net dipole moment of zero. In contrast, the water molecule (\(text{H}_2text{O}\)) also has polar bonds, but its bent molecular geometry causes the bond dipoles to add up instead of cancel, producing a substantial net dipole moment.
The Rule of Solubility: Why “Like Dissolves Like”
One of the most observable consequences of molecular polarity is solubility, which is governed by the principle that “like dissolves like.” This rule means that polar substances are generally soluble in polar solvents, and nonpolar substances dissolve well in nonpolar solvents. The basis for this behavior is the nature of intermolecular forces, the attractive forces that exist between molecules.
When a polar solute, such as table salt, is mixed with a polar solvent, like water, the partial charges on the solvent molecules can effectively attract and surround the charged particles of the solute, pulling them into solution. Nonpolar molecules, such as oils and fats, lack these permanent charges. When a nonpolar substance is placed in water, the strong attractions between the water molecules push the nonpolar molecules aside, causing them to separate and form distinct layers. Soap acts as a useful intermediary because its molecules are amphiphilic, possessing a long nonpolar hydrocarbon tail that interacts with grease and a polar head that interacts with water, allowing the two immiscible substances to mix.
Polarity’s Influence on Physical States and Properties
The attractive forces created by molecular polarity impact a substance’s physical properties, particularly its melting and boiling points. Because polar molecules have distinct positive and negative ends, they are subject to stronger intermolecular attractions, specifically dipole-dipole forces, which require more energy to overcome. Substances composed of polar molecules generally have significantly higher boiling points than nonpolar molecules of a similar size. For example, polar water boils at \(100^{circ}text{C}\), while nonpolar methane (\(text{CH}_4\)) boils at \(-161^{circ}text{C}\).
The strong cohesive forces between polar molecules also contribute to high surface tension, which is the ability of a liquid’s surface to resist external force. This property is involved in biological systems, where the tendency of polar molecules to interact and exclude nonpolar ones drives the formation of cell membranes. Cell membranes are structured as a lipid bilayer, with nonpolar tails tucked inside and polar heads facing the aqueous environment.