0.16 Phase equilibrium and intermolecular forces  (Page 4/7)

Figure 3

Figure 3 is an example of a complete phase diagram. This diagram shows for each temperature and pressure which phase or phases are present at equilibrium. Figure 3 is for water, but each substance has its own unique phase diagram, similar in appearance.

Observation 3: boiling points and intermolecular forces

Earlier in this study, we determined that the boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the pressure applied externally, e.g. by the atmosphere or by a piston trapping the liquid and gas in a cylinder. When the applied pressure is 1 atm, we refer to this as the “normal boiling point.” If we compare the normal boiling points of two liquids, the liquid with the higher normal boiling point clearly requires a higher temperature to reach a vapor pressure of 1 atm. From our work on dynamic equilibrium, we know that this higher temperature is required to provide sufficient kinetic energy for the molecules in the liquid to overcome stronger attractions between the molecules. Overcoming these attractions is necessary for a molecule to “escape” the liquid and join the vapor phase.

This line of reasoning means that, when we compare the normal boiling points of two liquids, we are also indirectly comparing the strengths of the intermolecular attractions in those two liquids. The liquid with a higher boiling point has stronger intermolecular attractions.

What determines the strength of these attractions? To find out, we can analyze experimental data for the boiling points of many liquids and look at the properties of the corresponding molecules. A useful set of compounds to look at are the covalent compounds formed by combining hydrogen with each of the elements in the “main group,” Groups IV to VII. For example, in the first row of the periodic table, these include CH ₄ , NH ₃ , H ₂ O, and HF. Table 1 gives the experimentally observed normal boiling points of the sixteen hydrides from Groups IV to VII in the first four rows of the periodic table.

At first glance, the values of the boiling points seem to be all over the place. Any patterns that might exist are not obvious. But there are patterns if we look at the data long enough, and those patterns can reveal to us what determines the intermolecular attractions. First, we can see that, for the compounds in each row of the periodic table, the compound with the lowest boiling point is from Group IV: CH ₄ , SiH ₄ , GeH ₄ , and SnH ₄ . Note that these are not the lowest four boiling points in Table 1. Rather, for each period, they are lowest boiling points of the compounds in each period. Notice also that the boiling points increase as we move down the table in Group IV, so that the heavier mass molecules have higher boiling points. This is easiest to see if we put them together on a chart in Figure 4.

These two observations suggest that we might find patterns if we put all of the sixteen hydrides on a chart together, sorted by the period each are in. The result is shown in Figure 5, which is just Figure 4 expanded to show all four groups in the data set of Table 1. With this chart, we see that the two patterns we described for Group IV work for Groups V to VII, but we also see three dramatic exceptions to those patterns in H ₂ O, NH ₃ , and HF.