Boiling Points

Liquids are held together by weak interactions between their molecules. To transition to the gas phase, a molecule must gain enough energy from heat or other sources to overcome these interactions.

Let's consider a closed container of a pure liquid at room temperature. Some of the liquid phase molecules at the surface gain enough energy to evaporate, and some of that vapor loses energy and condenses back to liquid. When the rates of evaporation and condensation are the same, the liquid and vapor reach equilibrium. At this point, the space above the liquid is saturated with vapor, and there is no net change in the amount of liquid and vapor in the container. The pressure exerted by the vapor is called the vapor pressure.

Now, consider an open container of liquid. Here the vapor can escape, keeping the area above the liquid from being saturated. Thus, evaporation could continue until there is no liquid left. As the temperature of a liquid increases, so does the evaporation rate, and therefore, so does the vapor pressure. If enough heat is added, molecules deeper in the liquid start vaporizing, which we see as bubbles of vapor forming.

This is called boiling, and it starts when the vapor pressure of the compound is equal to the atmospheric pressure. The temperature at which this occurs is the boiling point. Since each pure substance has different types and strengths of intermolecular interactions, different substances have different boiling points.

Now, let's take a look at how intermolecular interactions impact the boiling point. Many intermolecular interactions involve areas with uneven electron density called dipoles. All molecules have brief fluctuations in their electron distribution, so even nonpolar molecules can temporarily have dipoles. In response, opposing dipoles form in nearby molecules.

The attractive forces between temporary and induced dipoles are called London dispersion forces, and they are the dominant interactions between nonpolar molecules like hydrocarbons. London dispersion forces are stronger between large molecules with many electrons because those molecules hold their electrons weakly, making it easier for neighboring molecules to affect their electron distribution.

For example, a high molecular weight unbranched alkane has a higher boiling point than a lower molecular weight unbranched alkane. Additionally, an unbranched alkane has a higher boiling point than a branched alkane with a comparable molecular weight due to its higher surface area.

Molecules with polar bonds have permanent dipoles, enabling regular dipole-dipole interactions between them. Dipole-dipole interactions are stronger than London dispersion forces, so it usually takes more thermal energy to boil polar liquids than nonpolar liquids. For example, the boiling point of one chloropentane, a polar molecule, is 108 °C. Hexane, which is similar in surface area but lacks a dipole, has a lower boiling point.

Molecules with both a hydrogen covalently bound to a more electronegative atom and an electron-withdrawing atom with a lone pair of electrons are capable of hydrogen bonding. Hydrogen bonds are stronger than dipole-dipole interactions and London dispersion forces, so it takes even more thermal energy to overcome them. For example, propane and ethanol have similar molecular weights and surface areas. But the boiling point of propane is much lower than ethanol's because ethanol molecules can form hydrogen bonds with each other, while propane cannot.

In this lab, you will use the capillary method to determine the boiling point of two organic solvents.