Recall:

We were talking about:

Intermolecular Forces

Summary of intermolecular forces in pure substances:

Hydrogen bonding forces (H-bonding forces):

Only in molecules with an O- H bond, an N- H bond, or an F- H bond.

Strongest of the intermolecular forces between neutral molecules.

Approximately 10% as strong as an actual covalent single bond.

Dipole-dipole forces:

Only in polar molecules

Depends on the size of the molecular dipole moment.

Approximately 5% as strong as H-bonding forces.

Dispersion forces:

Present in all molecules.

Roughly proportional to the square of the "volume" of the molecule.

H₂, HCl, Cl₂, H₂S, C₆H₆, CH₃OH, CO₂, HCN, glycerine

Vapor Pressure and Boiling Points

We can use our knowledge of intermolecular forces to compare vapor pressures and boiling points of liquids.

Vapor Pressure

Vapor pressure is a measure of the ability of molecules to escape from the surface of a liquid (or a solid, for that matter). Vapor pressure is sometimes referred to as "escaping tendency."

We know that there is a competition between intermolecular forces (attracting molecules to each other) and the kinetic energy of the molecules (which is trying to separate them).

We also know that the average kinetic energy of the molecules is proportional to the Kelvin temperature, T. However, not all molecules have the average kinetic energy. There is a distribution of kinetic energies with some molecules having a kinetic energy higher than the average and some having an energy lower than the average.

When the temperature is increased the fraction of molecules with sufficient kinetic energy to overcome the intermolecular forces increases so that the vapor pressure increases.

If the temperature is decreased the vapor pressure decreases because fewer molecules have enough kinetic energy to escape the intermolecular forces.

We can compare vapor pressures of two different substances at the same temperature.

At a given temperature, the substance with the stronger intermolecular forces will have the lower vapor pressure.

CH₄, NH₃, H₂O

H₂O, H₂S, H₂Se

A liquid boils when its vapor pressure reaches one atmosphere.

We can use our knowledge of intermolecular forces to predict the relative boiling points of compounds.

Since compounds with large intermolecular forces have lower vapor pressures, we predict that one has to go to higher temperature to make them boil.

So we can use the following principles to predict relative boiling points:

Stronger intermolecular forces mean higher boiling points.

Other things being equal:

Polar molecules boil higher than nonpolar molecules.

HCl                       - 82^oC
Ar                        - 186^oC

CH₄                     - 162^oC
NH₃                      - 34^oC
H₂O                       100^oC

He                         4.22 K
Ar                          87.5 K

CH₃(CH₂)₃CH₃      36^oC
C(CH₃)₄                 9.5^oC

Intermolecular Forces in Mixtures

In mixtures we get a variety of new combinations of molecular interactions. For example in water solutions of ionic compounds there are interactions of the dipole moment of the water molecules with the charge on the ions to produce "solvated ions."

Sometimes the solvation energy is strong enough to carry water molecules into the crystals when the ionic compounds are crystallized out. This is called "water of crystallization," and gives rise to formulas such as BaCl₂×2H₂O.

In nonionic solutions and in gas phase mixtures there are other combinations of intermolecular forces. For example, the interaction between a polar molecule and a nonpolar molecule is stronger than dispersion forces and weaker than dipole-dipole forces.

Intermolecular forces give us an indication of the solubility of a compound in another compound. The principle is that "like dissolves like." That is, nonpolar solvents tend to dissolve other nonpolar compounds but not, for example, hydrogen-bonded compounds. Polar solvents tend to dissolve other polar compounds, but not nonpolar compounds or hydrogen bonded compounds.
 

Superheated and Supercooled Liquids

Sometimes, if the sample and container are very clean and there are no dust particles in the liquid, you can heat a liquid above it's boiling point without it boiling. This is called a superheated liquid. A superheated liquid is not at equilibrium - the liquid is unstable. If you stir the liquid or drop in a grain of salt or something similar the system will release all of its excess energy all at once in a flash of vapor. (It might look like a small explosion.)

In the laboratory, when boiling liquids, say in distillation, we add "boiling stones" to the liquid. The boiling stones provide sites for the nucleation of bubbles and keep the liquid from overheating.

By the same token, liquids can be cooled below their freezing point without freezing - if it is carefully done and the system is very pure. This results in what is called a supercooled liquid. A supercooled liquid is not at equilibrium and will freeze immediately if the system is stirred or a seed crystal of the solid material is added to the system.
 

Phase Diagrams

Many of the properties of a material can be summarized on a figure called a phase diagram. A phase diagram is a set of curves drawn in relation to a set of coordinate axes where pressure, p, is plotted on the vertical axis and temperature, T, is plotted on the horizontal axis.

A typical phase diagram might look like this:
 
 

The lines separate the graph into regions where the compound exists as a solid, a liquid or a gas.

We can identify the melting and boiling points on the graph by locating where one atmosphere appears on the vertical axis.

The line which separates the solid region from the gas region is called the sublimation curve.

The line which separates the solid region from the liquid region is called the melting curve.

And the line which separates the liquid region from the gas region is called the vaporization curve. This curve is also called the vapor pressure curve because it is a plot of the vapor pressure of the liquid versus temperature.

The vaporization curve has an end at a point called the "critical point." The critical point is defined by a critical pressure, p_c, and a critical temperature, T_c. A gas can be liquefied at a temperature below the T_c, if we go to high enough pressure, but the gas can not be liquefied at any pressure if the temperature is above T_c.

The point where the three phase lines meet is called the "triple point" because at this point, and only this point, the three phases, solid, liquid, and gas, can exist all at the same time in equilibrium with each other.

You will notice that on our phase diagram the temperature of the triple point is lower than the melting point. This is the usual case. However, for water the triple point is 273.16 K and the melting point is 273.15 K. The phase diagram for water would look qualitatively different, like,

The negative slope of the solid/liquid line has an enormous impact on life on our planet. (I have exaggerated the negative slope on this diagram. On this scale it would look like the line went straight up.)

You are probably aware that solid CO₂ does not melt at one atmosphere pressure. The phase diagram for CO₂ would look something like,

Crossing a line requires a gain or loss of energy. Crossing the line from solid to liquid requires that the system absorb energy, which we called DH_fus, and crossing from liquid to gas requires an energy which we called DH_vap. (chapter 6)

The slopes of the lines on a phase diagram can be related to the DH of the phase transition involved. One example, called the Clausius-Clapeyron equation is given in your text on page 604.