Intermolecular forces are the attractions between individual molecules that dictate how a substance behaves in the physical world. These forces are distinct from the strong covalent or ionic bonds that hold atoms together within a molecule. They govern many observable properties, such as a compound’s melting point, boiling point, and solubility. Understanding these attractions is fundamental to predicting the bulk properties of any substance.
Types of Intermolecular Forces
The forces acting between molecules are categorized into three main types, differing primarily in strength. The weakest are London Dispersion Forces (LDF), which result from the momentary movement of electrons within a molecule. These temporary shifts create instantaneous dipoles that induce similar dipoles in neighboring molecules, and LDFs are present in all substances.
Dipole-dipole forces occur only in molecules that possess a permanent separation of charge. This electrical imbalance allows the slightly positive end of one molecule to attract the slightly negative end of a neighbor. The strongest attraction is hydrogen bonding, a special, amplified form of dipole-dipole interaction. Hydrogen bonding requires a hydrogen atom covalently linked to a highly electronegative atom like nitrogen, oxygen, or fluorine.
How Molecular Shape Determines Polarity
For dipole-dipole forces to be active, a molecule must be polar, meaning it has a permanent dipole moment. Polarity begins with the bonds themselves: a covalent bond between atoms with different electronegativities is polar due to unequal electron sharing. The more electronegative atom pulls the shared electrons closer, acquiring a partial negative charge, while the other atom gains a partial positive charge.
The overall polarity of the molecule depends on its three-dimensional geometry. If a molecule has polar bonds but a highly symmetrical shape, the individual bond dipoles cancel each other out. This results in a nonpolar molecule, such as carbon dioxide or boron trifluoride, despite the presence of polar bonds. For a molecule to be polar and exhibit a permanent dipole moment, it must possess an asymmetrical structure that prevents this cancellation.
Analyzing Nitrogen Trifluoride
Nitrogen trifluoride (\(\text{NF}_3\)) is composed of one nitrogen atom bonded to three fluorine atoms. The individual N-F bonds are polar because fluorine is much more electronegative than nitrogen, causing electrons to be pulled toward the fluorine atoms. Therefore, the molecule contains three distinct bond dipoles.
The nitrogen atom in \(\text{NF}_3\) has four electron domains: three bonding pairs and one non-bonding lone pair. This arrangement results in a trigonal pyramidal molecular shape. The lone pair occupies space on the central nitrogen atom, pushing the three fluorine atoms downward and creating an asymmetrical structure.
This trigonal pyramidal geometry prevents the three N-F bond dipoles from canceling one another out. The lone pair on the nitrogen atom also contributes a dipole moment directed away from the fluorine atoms. The combination of the uncancelled bond dipoles and the lone pair’s influence gives \(\text{NF}_3\) a measurable net dipole moment. Since \(\text{NF}_3\) possesses a permanent dipole, it exhibits dipole-dipole forces between its molecules, in addition to London Dispersion Forces.
Consequences of Dipole Dipole Interaction
The presence of dipole-dipole forces means that \(\text{NF}_3\) molecules are attracted to one another with a greater force than if only London Dispersion Forces were active. This stronger attraction translates into observable differences in physical properties. Substances with stronger intermolecular forces require more energy to overcome those attractions and transition from a liquid to a gas.
Consequently, \(\text{NF}_3\) has a higher boiling point compared to a nonpolar molecule of similar size that relies only on LDFs. For example, when compared to boron trifluoride (\(\text{BF}_3\)), which is nonpolar due to its symmetrical shape, \(\text{NF}_3\)‘s dipole-dipole forces allow it to remain a liquid at a higher temperature. The enhanced intermolecular attraction also influences solubility, as polar \(\text{NF}_3\) will be more attracted to other polar solvents.