Intermolecular forces (IMFs) are the attractive or repulsive forces that exist between neighboring molecules, acting like a temporary “glue” that holds matter together in condensed phases, such as liquids and solids. These forces are significantly weaker than intramolecular bonds, like covalent or ionic bonds, which hold atoms together within a single molecule. Dipole-dipole forces are one specific type of IMF, arising from the electrical attraction between certain molecules. Understanding this specific force is fundamental to explaining many physical properties of substances.
Molecular Requirements for Dipole Forces
The existence of a dipole-dipole force depends entirely on a molecule possessing a permanent dipole. This permanent electrical imbalance is a consequence of the unequal sharing of electrons between atoms, a process driven by differences in electronegativity. Electronegativity is an atom’s ability to attract shared electrons; when bonded atoms have a notable difference in this property, electrons are pulled closer to the more electronegative atom.
This uneven electron distribution creates a separation of charge, resulting in one end of the molecule having a partial negative charge (\(\delta-\)) and the other end having a partial positive charge (\(\delta+\)). The molecule is then considered polar, acting as a tiny, permanent electric dipole. For a molecule to be polar, the individual bond dipoles must not cancel each other out due to molecular symmetry.
For example, in hydrogen chloride (HCl), the more electronegative chlorine atom pulls electrons closer, becoming the partial negative pole and leaving the hydrogen atom as the partial positive pole. This molecular asymmetry creates a net dipole moment, which is the necessary prerequisite for the interaction. Conversely, carbon dioxide (CO\(_2\)) has polar bonds, but its linear shape causes the dipoles to perfectly cancel, resulting in a nonpolar molecule that cannot participate in dipole-dipole interactions.
How Dipole-Dipole Forces Interact
Once a molecule possesses a permanent dipole, it can engage in a dipole-dipole interaction with another polar molecule. This interaction is fundamentally an electrostatic attraction, where the partial positive pole of one molecule is drawn toward the partial negative pole of a neighbor. The oppositely charged ends are attracted to each other, similar to the poles of two small magnets.
This attraction is highly directional and tends to cause polar molecules to align themselves to maximize attractive forces. In condensed phases, molecules will shift positions to ensure the positive end of one dipole faces the negative end of an adjacent dipole. This alignment also minimizes the repulsive forces that occur if two like poles face one another.
The strength of this force is inversely related to the distance between molecules, meaning it is only effective over very short ranges. Because the charges involved are only partial charges, dipole-dipole forces are much weaker than the full electrical charge attraction seen in ionic bonds. However, the permanent nature of the dipoles makes this attraction a significant factor in how polar substances behave.
Ranking Intermolecular Force Strength
Dipole-dipole forces fall into an intermediate range of strength compared to the two other main types of intermolecular forces. They are generally stronger than London Dispersion Forces (LDF), which are temporary attractions existing in all molecules. For molecules of comparable size, the presence of a permanent dipole makes the total intermolecular attraction stronger than if only LDF were present.
The magnitude of the dipole moment, a measure of charge separation, determines the ultimate strength of the attraction for a given molecule. A molecule with a greater charge difference between its poles will exhibit a stronger dipole-dipole force.
This force is significantly less powerful than hydrogen bonding. Hydrogen bonding is a specialized interaction that occurs only when hydrogen is bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine, creating an extremely strong partial positive charge. This hierarchy places dipole-dipole forces above LDF but below hydrogen bonding.
How Dipole Forces Affect Physical Properties
The presence of dipole-dipole forces has direct consequences on a substance’s physical properties. Because these attractions hold molecules together more tightly than LDF alone, more external energy is required to separate them. This requirement translates directly into higher boiling points and melting points for polar compounds compared to nonpolar compounds of similar mass.
For instance, two molecules with nearly identical molecular weights, one polar and one nonpolar, will show a distinct difference in their phase change temperatures. The polar molecule requires additional energy to break the dipole-dipole forces.
Dipole-dipole interactions also play a major role in solubility, following the principle that “like dissolves like.” Polar substances tend to dissolve well in other polar solvents because the new attractive forces formed between the solute and solvent are comparable in strength to the original dipole-dipole forces.