A dipole in chemistry describes a separation of electrical charge within a molecule, resulting in a positive end and a negative end. This charge unevenness means that the molecule, or a part of it, has two “poles” with opposite electrical properties. This separation occurs when electrons are not shared equally between atoms in a chemical bond. The presence of a dipole gives a molecule an overall electrical asymmetry, which is a fundamental property influencing how it interacts with other matter.
The Underlying Cause of Polarity
The mechanism that drives the formation of a dipole is the difference in a property called electronegativity between bonded atoms. Electronegativity is an atom’s inherent ability to attract a shared pair of electrons toward itself within a chemical bond. When two atoms with significantly different electronegativity values form a covalent bond, the shared electrons spend more time orbiting the more attractive atom. This unequal sharing creates an imbalance in electron density.
The atom that attracts the electrons more strongly develops a slight excess of negative charge, represented by the Greek letter delta negative (\(\delta-\)). Conversely, the less attractive atom develops a slight positive charge, denoted as delta positive (\(\delta+\)). This partial charge separation constitutes the bond dipole. A bond between identical atoms, like in \(\text{H}_2\), is nonpolar, but a bond between hydrogen and chlorine (\(\text{HCl}\)) has a large electronegativity difference, making it a polar bond.
Distinguishing Between Bond and Molecular Dipoles
It is important to differentiate between the localized polarity within a single bond and the overall polarity of an entire molecule. A bond dipole refers to the charge separation that exists between just two bonded atoms, a concept determined solely by the electronegativity difference between those two atoms. Every polar covalent bond possesses a bond dipole, which is a vector quantity having both magnitude and direction.
The molecular dipole, on the other hand, is the net, overall polarity of the whole molecule. It is the result of the vector sum of all individual bond dipoles within the molecule, and it is heavily influenced by the molecule’s three-dimensional shape. If the bond dipoles are arranged symmetrically, they can effectively cancel each other out, leading to a molecular dipole of zero.
For example, the linear carbon dioxide (\(\text{CO}_2\)) molecule has two polar carbon-oxygen bonds, but because they pull electrons in opposite and equal directions, the molecule as a whole is nonpolar. In contrast, a bent molecule like water (\(\text{H}_2\text{O}\)) has two highly polar bonds whose individual dipoles do not cancel due to the angular geometry. This structural asymmetry results in a net molecular dipole moment, making water a highly polar molecule.
How Dipoles Are Measured and Represented
Chemists quantify the magnitude of this charge separation using a value called the dipole moment, symbolized by the Greek letter mu (\(\mu\)). The dipole moment is mathematically calculated as the product of the magnitude of the partial charge (\(Q\)) and the distance (\(r\)) separating the charges. The standard unit for measuring the dipole moment in chemistry is the Debye (D).
Visually, the polarity of a bond or a molecule is represented by a dipole arrow. The arrow is drawn parallel to the bond axis, with a cross at the positive end to denote the partial positive charge (\(\delta+\)). The arrowhead points toward the more electronegative atom or the region of higher electron density, which is the partial negative charge (\(\delta-\)). This arrow is a vector, conveying both the direction and the magnitude of the electron pull.
How Dipoles Influence Chemical Behavior
The existence of a molecular dipole has profound consequences for a substance’s physical and chemical behavior. Polar molecules align themselves so that the positive end of one molecule is attracted to the negative end of a neighboring molecule, an interaction known as a dipole-dipole force. These forces are a type of intermolecular force, which are the attractions that exist between separate molecules.
The strength of these attractions directly influences physical properties. Polar substances generally have higher melting and boiling points than nonpolar substances of comparable size because more energy is required to overcome these cohesive forces. A particularly strong form of dipole-dipole interaction is hydrogen bonding, which occurs when hydrogen is bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine.
Furthermore, the principle of “like dissolves like” governs solubility. Polar molecules are typically soluble in polar solvents, such as water, because the strong dipole-dipole attractions stabilize the mixture. Conversely, polar molecules do not readily dissolve in nonpolar solvents, as the attractions between the polar molecules are much stronger than the weak forces that would form between the polar and nonpolar molecules.