The dipole force is an attractive interaction that occurs between molecules possessing permanent electrical asymmetry. It is one of the three main types of intermolecular forces, which are the attractions and repulsions between neighboring particles. Essentially, it is an electrostatic pull between the positive end of one molecule and the negative end of an adjacent molecule. This force dictates many physical properties of a substance, such as its boiling point and solubility.
How Molecular Polarity Arises
Molecular polarity, the prerequisite for a dipole force, originates from the unequal sharing of electrons within a covalent bond. This unequal sharing is quantified by electronegativity, which is an atom’s ability to attract a shared electron pair toward itself. When two atoms with different electronegativities bond, the more electronegative atom pulls the electron cloud closer, acquiring a partial negative charge (delta minus).
The less electronegative atom is left with a partial positive charge (delta plus), creating a separation of charge within the bond. For example, in Hydrogen Chloride (HCl), chlorine is more electronegative than hydrogen, causing electrons to spend more time near the chlorine. This charge separation establishes a bond dipole moment, a vector quantity with both magnitude and direction.
A molecule develops an overall polarity, or a net dipole moment, based on the combined effect of all its bond dipoles and its three-dimensional geometry. Even if a molecule contains polar bonds, a symmetrical shape can cause the individual dipoles to cancel out, resulting in a nonpolar molecule. If the polar bonds are arranged asymmetrically, like in a bent water molecule, the molecule retains a permanent separation of charge, making it polar.
The Dipole-Dipole Attraction
The dipole-dipole attraction is the direct interaction between neighboring polar molecules. This attraction occurs when the partial positive pole (delta plus) of one molecule aligns with the partial negative pole (delta minus) of a second molecule. This electrostatic alignment maximizes the attraction between opposite charges while minimizing the repulsion between like charges.
The strength of this force is moderate among intermolecular forces, being significantly weaker than the forces within a molecule, such as ionic or covalent bonds. Dipole forces are much stronger than London Dispersion forces, which are temporary attractions present in all molecules. The energy required to break these attractions ranges from 5 to 20 kilojoules per mole (kJ/mol).
The magnitude of the dipole-dipole force is directly related to the strength of the molecule’s permanent dipole moment. The interaction is highly dependent on the distance between the molecules, weakening rapidly as they move farther apart. For molecules of similar size, the presence of a permanent dipole moment makes the dipole-dipole interaction the dominant attractive force compared to nonpolar interactions.
Influence on Substance Properties
Dipole-dipole forces affect the physical characteristics of a substance because these attractions must be overcome to change a substance’s state. For molecules of comparable size, those that exhibit dipole-dipole forces require more energy to separate than those that rely only on weaker dispersion forces. This increased energy requirement results in higher melting points and boiling points.
The influence of these forces is apparent in solubility, often summarized by the phrase “like dissolves like.” Polar substances dissolve readily in polar solvents, such as water, because the strong dipole-dipole forces between the solute and solvent molecules replace the attractions between the solute molecules. Conversely, polar substances do not dissolve well in nonpolar liquids because the energy gained from weak attractions is not enough to break the strong attractions within the polar substance.
A strong form of the dipole-dipole interaction is known as hydrogen bonding. This occurs when a hydrogen atom is bonded directly to a highly electronegative atom like nitrogen, oxygen, or fluorine. This interaction, exemplified by the structure of water, raises the boiling point and is responsible for many of the unusual properties of water.