Even though NaCl does not dissolve in hexane, \(\mathrm{C}_{6} \mathrm{H}_{14}\), if we imagine the dissolving process for this system, there is one step that would take less energy than the corresponding step for \(\mathrm{NaCl}\) dissolving in water. Which step is it? Why would it require less energy?

The process of dissolving a substance typically involves a crucial step: the separation of solvent particles. To make room for solute particles, such as ions of an ionic compound, we must disrupt the established intermolecular forces within the solvent.

In the case of water, these forces are manifested as hydrogen bonds, which are quite strong due to the polarity of water molecules. These bonds create a tight-knit network that requires a significant amount of energy to break apart. Conversely, a solvent like hexane exhibits much weaker forces, known as London dispersion forces. These are induced momentarily due to fluctuating electron densities and result in a weaker hold between molecules.

When discussing solubility, this is an essential factor. The weaker the force holding the solvent molecules together, the less energy it requires to separate them, making room for the solute. Hence, the reason why separating hexane molecules demands less energy compared to separating water molecules—leading to the key insight that while not all steps of dissolution are equal, the energy demand for separation of solvent particles can vary significantly between solvents.

Intermolecular forces are the attractive or repulsive forces between molecules, including ionic or covalent substances. These forces play a fundamental role in the physical properties of substances, particularly solubility. The major types of intermolecular forces include London dispersion forces, dipole-dipole interactions, and hydrogen bonds, each varying in strength and impact on a molecule's behavior in different environments.

• London dispersion forces are the weakest and present in all molecules, becoming significant in nonpolar substances like hexane.
• Dipole-dipole interactions occur between polar molecules and are stronger than London forces.
• Hydrogen bonds are a special case of dipole-dipole interactions where hydrogen is involved, resulting in a remarkably robust bond—as seen in water.

These forces determine how much energy it takes to break apart molecules in a solvent to accommodate a solute. Hence, a solvent like water with strong hydrogen bonds will require more energy to disrupt than hexane, which is governed by the much weaker London dispersion forces.

After the initial separation of solvent particles, the interaction between the solute and solvent comes into play. With ionic compounds, this interaction is particularly critical and is known as ion-dipole interaction. It occurs between the charged ions of the solute and the polar molecules of the solvent.

When a salt like NaCl is introduced to water, the polar water molecules align themselves according to the charge distribution of the ions. The positive end of the water molecule (hydrogen side) is attracted to the chloride ion (Cl-), while the negative end (oxygen side) is attracted to the sodium ion (Na+). These strong attractions help stabilize the ions in solution, contributing to the solubility of NaCl in water.

In contrast, in a nonpolar solvent like hexane, there's no significant charge separation within the molecules, making ion-dipole interactions essentially nonexistent. This lack of interaction means that ionic compounds like NaCl have virtually no tendency to dissolve in nonpolar solvents such as hexane, as the energy released from such weak interactions is insufficient to make the dissolution process favorable.