So far, we've discussed molecules with permanent dipoles. The differences in electronegativity between the bonding atoms create local areas of net positive or negative charge that remain in place because the bonding electrons do not shift much. But it can also happen that, from time to time, that all the electrons in a molecule just coincidentally wind up on the same side of the molecule, even if it is non-polar, like diatomic O2. Such temporary forces have a net effect at each instant on the behavior of the substance.
These momentary dipoles can be coincidental (dispersion), but they can also be induced by the presence of a strongly repulsive or attractive ion or dipole molecule nearby (induced London Forces). The temporary dipoles are constantly changing as the forces that give rise to them change (molecules move about) and are weak compared to other intermolecular forces, but in a large vat of liquid, they do actually have a net result....the liquid tends to resist just flying apart and becoming a gas.
Large electron clouds tend to form dispersion areas easily. This tendency is called polarization, and is responsible for the attraction of halogens into diatomic molecules, such as I2.
In summary, the temporary intermolecular forces are:
Molecules are pretty easy to contemplate: they are discrete units, with definite and consistent composition. They can be separated and still keep their chemical identity as they change phase from solid to liquid to gas. As gases, the molecules move in isolation from one another, but in liquid form, the molecules are close enough to be slightly attracted by intermolecular forces, and as solids, the forces are strong enough to maintain spatial relationships.
Not all solids are made of discrete molecules, however, and this affects the types of physical and chemical change these materials sustain.
Molecules and atoms in the liquid phase have individual kinetic energy large enough to keep them moving, constantly breaking and creating relationships with the molecules around them. Collisions will give some of the molecules enough energy to break free of even these tenuous relationships and enter the gas phase, so the surface of a liquid is a place of constant recombination, as some molecules become gaseous and others are captured by intermolecular forces and trapped in the liquid. The energy gained or lost in this transition is the enthalpy of vaporization.
The rate at which molecules become gas is not necessary the rate at which molecules become liquid across the same surface. The two depend on pressure and temperature and the nature of the intermolecular forces. At any given pressure, there is a temperature at which the two rates will be the same, and this is the equilibrium vapor pressure. If the pressure remains the same and the temperature rises, the substance will evaporate as more molecules become gas than become liquid. If the temperature drops for the same pressure, the substance condenses. The temperature, enthalpy, and pressure are all related, and the relationship was simultaneously discovered by a German physicist named Clausius, and a French chemist, Clapeyron, so the equation is the Clausius-Clapeyron equation:
The symbol "ln" stands for the natural logarithm function, a logarithm that uses the transcendental number e as its base instead of 10. If ex = y, then ln ex = x = ln y. ex is a special curve: its slope is always equal to its value. This makes it useful for plotting physical changes that involve rates or ratios, like the ratio of P to T. In this case, the equation means that
P = e-ΔH/RT; the C is an offset depending on nature of the compound in question. R is our old friend, the gas constant.
Review the discussion of macroscopic and microscopic explanations for vapor pressure behavior at Purdue's chemistry department website.
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