At \(35^{\circ} \mathrm{C}\) the vapor pressure of acetone, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO},\) is 360 torr, and that of chloroform, \(\mathrm{CHCl}_{3}\), is 300 torr. Acetone and chloroform can form very weak hydrogen bonds between one another as follows: A solution composed of an equal number of moles of acetone and chloroform has a vapor pressure of 250 torr at \(35^{\circ} \mathrm{C}\). (a) What would be the vapor pressure of the solution if it exhibited ideal behavior? (b) Use the existence of hydrogen bonds between acetone and chloroform molecules to explain the deviation from ideal behavior. (c) Based on the behavior of the solution, predict whether the mixing of acetone and chloroform is an exothermic \(\left(\Delta H_{\text {soln }}<0\right)\) or endothermic \(\left(\Delta H_{\text {soln }}>0\right)\) process.

Vapor pressure is a measure of a liquid's tendency to evaporate into a gas. It's determined by how many molecules escape from the liquid to create a saturated vapor above the liquid at a given temperature. The higher the vapor pressure, the more volatile the substance, indicating that it evaporates readily.

When considering an ideal solution, Raoult's Law provides that the total vapor pressure is the sum of the partial pressures of each component. The partial pressure is the product of the mole fraction of that component and its pure vapor pressure. In the context of our textbook example, for a mixture of acetone and chloroform in equal molar amounts, the expected vapor pressure would be calculated by adding the partial pressures, which are themselves half of the pure component vapor pressures. However, the actual pressure is lower due to deviations from ideality, which brings us to the role of intermolecular forces, specifically hydrogen bonding.

Hydrogen bonding is a type of dipole-dipole interaction that occurs when a hydrogen atom, which is bonded to a highly electronegative atom such as nitrogen, oxygen or fluorine, is attracted to another electronegative atom. In the case of acetone and chloroform, although the hydrogen bonding is weak, it still impacts the behavior of the mixture.

These weak hydrogen bonds alter vapor pressure by stabilizing the liquid phase, effectively reducing the number of molecules that can escape into the vapor phase. The stronger the intermolecular forces within a liquid, the lower its vapor pressure will be at a given temperature. Thus, in our textbook exercise, the actual vapor pressure of 250 torr, compared to the expected 330 torr for an ideal solution, suggests that additional intermolecular forces—in this case, hydrogen bonds—are at play, holding the molecules of acetone and chloroform together more tightly than if they were in separate pure substances.

The enthalpy change of mixing, \( \Delta H_{\text{mix}} \), refers to the heat absorbed or released when substances are mixed. If heat is released during mixing, the process is exothermic \( \Delta H_{\text{mix}} < 0 \), indicating that the energy required to break original bonds is less than the energy released from forming new ones. When heat is absorbed, the process is endothermic \( \Delta H_{\text{mix}} > 0 \), meaning that it takes more energy to break the initial intermolecular forces than is released upon forming the new bonds.

In our textbook exercise, the mixture of acetone and chloroform has a lower vapor pressure than what Raoult's Law predicts for an ideal solution, demonstrating stronger intermolecular forces due to hydrogen bonding. This is indicative of an exothermic reaction—mixing acetone and chloroform releases heat, resulting in a negative enthalpy change of mixing. Essentially, the energy given off as the weak hydrogen bonds form is greater than what was used to separate the molecules initially.