IMFs in Solids and Liquids

As we saw earlier, one of the most important distinctions between gases and the condensed phases is that in solids and liquids, particles are close enough to each other to have sustained intermolecular interactions. We will now discuss the role of intermolecular forces - as well as chemical bonds, in some cases - in phase behavior. We will talk about the way energy is absorbed and released when these attractive forces form and break, as well as how they affect important physical properties of solids and liquids.

Energy and IMFs

It is the ionic bonds of NaCl that hold solid salt crystals together, and it is these ionic bonds that are partially broken when salt melts (yes, salt can melt! It happens around 800 °C). In water, it is hydrogen bonds that break when water is vaporized into a gas.

Because these are all forces of attraction, they require energy to break them. This is why heat is always needed to melt or vaporize a substance: the heat energy is "used" to pull the particles (ions, atoms, or molecules) apart from each other, breaking their bonds/IMFs. You can observe this on the macroscopic scale: if you take two magnets that are in contact with each other, energy is needed to separate them.

On the other hand, when two magnets are allowed to come together, they release energy in the form of a loud "clack." This also occurs when bonds and intermolecular forces form (such as when a gas condenses or a liquid freezes). Energy is released in the form of heat in those cases.

Endothermic and Exothermic

All processes that break or form chemical bonds or intermolecular forces can be classified as either endothermic or exothermic. This includes phase changes, as well as all chemical reactions.

• Some processes only involve breaking bonds/IMFs. Examples of these include melting, vaporization, and sublimation, as well as a small number of chemical reactions. These processes are always endothermic, which means they absorb heat. Examples include the melting of water and the sublimation of carbon dioxide (dry ice).

• Other processes only involve forming bonds/IMFs. These include the reverses of the above processes: freezing, condensation, and deposition (the reverse of sublimation). These processes are always exothermic, meaning they release heat. For whatever reason, students tend to have a more difficult time internalizing this bullet point than the one above. It may be useful to think of these processes in pairs and remember that energy is always conserved: melting is a process that absorbs heat, so reversing that process must release heat.

• Most chemical reactions involve both the breaking and forming of bonds/IMFs. These processes can be endothermic or exothermic, depending on the balance of these two processes. A single reaction might involve breaking some bonds (which requires energy) and forming some others (releasing energy). If the amount of energy released is greater than the energy absorbed, the reaction will be exothermic; in the reverse case, it will be endothermic.

IMFs and Physical Properties

First, as you can see, substances with stronger forces between their particles require higher temperature to melt them. This makes sense: higher temperature means more kinetic energy, because more energy is needed to overcome a stronger force.

Second, you can probably see from this table why we classify "intermolecular forces" and "chemical bonds" separately. Even though they both lie on a continuum of attractive forces, and that is how we are treating them here, there is a pretty clear gulf between substances held together by IMFs, which mostly melt well below room temperature, and substances held together by bonds, which mostly melt well above it.

It's worth emphasizing that there are other factors besides the type of IMF/bonding in a substance that go into determining its physical properties, including atomic/molecular size among many others. Just as one example: mercury, which has metallic bonding, melts at -38 °C, much lower than most metals.

Melting and Boiling Points

In general, though, the pattern is quite clear: substance with stronger attractive forces (IMFs or bonds) melt and boil at higher temperatures. In addition, there are many other properties that are influenced by IMF strength.

Viscosity

Viscosity is resistance to flow. Liquids with low viscosity, like water, flow easily. Liquids with high viscosity, like motor oil or syrup, flow only slowly. Generally, the more strongly the molecules are attracted to one another, the more they resist flowing past one another, and the higher the viscosity.

However, there are other important factors that go into determining viscosity, particularly the size of molecules and even their shape. Liquids made of larger molecules are generally more viscous, especially when those molecules are long and flexible, able to get tangled up with each other. For these reasons, olive oil (which is made of very large, long, flexible molecules) is nearly 100 times more viscous than water, despite having mostly dispersion forces.

Temperature has an important effect on viscosity. In general, liquids become significantly les viscous as temperatures drop; water is six times less viscous at room temperature than at its boiling point, and a drop of just 10 °C in the temperature of olive oil can nearly double its viscosity.

Surface Tension

Surface tension is the resistance of the surface of a liquid to being broken. It is caused by the fact that the molecules on the surface of a liquid are not bonded to as many neighboring molecules as those beneath the surface. As with viscosity, the greater the attraction between molecules, the higher the surface tension.

An object will break the surface of a liquid only if it exerts a force (due to gravity, momentum, etc.) that is greater than the force of the intermolecular bonds that must be broken.

Vapor Pressure

At any given temperature the molecules in a liquid have a broad range of kinetic energies – some are very high, some are very low, but most are close to the average. To escape from the liquid, a surface molecule must have sufficient kinetic energy to overcome the intermolecular forces holding it to its neighbors. The higher the temperature, the greater the fraction of surface molecules that will have this amount of energy and the greater the rate of evaporation. Further, at any given temperature, in a liquid whose molecules are held together by weak intermolecular forces, a relatively larger fraction of the surface molecules will have sufficient kinetic energy to overcome those forces than in a liquid whose molecules are held together by strong intermolecular forces.

The molecules now in the gas phase above the liquid are in constant motion and strike each other, the walls of the container, and the surface of the liquid. Those that strike the surface of the liquid often “stick,” attracted by the intermolecular forces that held them in the liquid in the first place, and recondense into the liquid phase. The rate at which this happens depends mostly on how many such collisions take place, which in turn depends on the number of gas molecules there are above the liquid. This number also determines the pressure the gas exerts, hence, the rate of condensation increases as the pressure increases.

Perhaps surprisingly, the vapor pressure of a substance does not depend on the size of the container or the amount of liquid in it ... so long as there is enough liquid that equilibrium is reached before all the liquid evaporates, the pressure exerted by the vapor phase is always the same. However, vapor pressure does depend significantly on temperature. For example, at 1 °C water has a vapor pressure of just 5 torr, whereas at 99 °C it is 758 torr.

As mentioned previously, substances with weak IMFs (especially smaller molecules) tend to have higher vapor pressures because molecules vaporize more easily.