0.2 Phase equilibrium and intermolecular interactions  (Page 6/7)

is still not a complete phase diagram, because we have not includedthe combinations of temperature and pressure at which solid and liquid are at equilibrium. As a starting point for theseobservations, we look more carefully at the conditions near 0°C. Very careful measurements reveal that the solid-gasline and the liquid-gas line intersect in where the temperature is 0.01°C. Under these conditions, we observe inside thecontainer that solid, liquid, and gas are all three at equilibrium inside the container. As such, this unique temperature-pressurecombination is called the triple point . At this point, the liquid and the solid have the same vapor pressure,so all three phases can be at equilibrium. If we raise the applied pressure slightly above the triple point, the vapor must disappear.We can observe that, by very slightly varying the temperature, the solid and liquid remain in equilibrium. We can further observe thatthe temperature at which the solid and liquid are in equilibrium varies almost imperceptibly as we increase the pressure. If weinclude the solid-liquid equilibrium conditions on the previous phase diagram, we get this , where the solid-liquid line is very nearly vertical.

Each substance has its own unique phase diagram, corresponding to the diagram in for water.

Observation 4: dynamic equilibrium

There are several questions raised by our observations of phase equilibrium and vapor pressure. The first wewill consider is why the pressure of a vapor in equilibrium with its liquid does not depend on the volume of the container intowhich the liquid evaporates, or on the amount of liquid in the container, or on the amount of vapor in the container. Why do weget the same pressure for the same temperature, regardless of other conditions? To address this question, we need to understand thecoexistence of vapor and liquid in equilibrium. How is this equilibrium achieved?

To approach these questions, let us look again at the situation in . We begin with a container with a fixed volume containing some liquid,and equilibrium is achieved at the vapor pressure of the liquid at the fixed temperature given. When we adjust the volume to a largerfixed volume, the pressure adjusts to equilibrium at exactly the same vapor pressure.

Clearly, there are more molecules in the vapor after the volume is increased and equilibrium is reestablished,because the vapor exerts the same pressure in a larger container at the same temperature. Also clearly, more liquid must haveevaporated to achieve this equilibrium. A very interesting question to pose here is how the liquid responded to the increase in volume, which presumably only affected the space in which the gas moleculesmove. How did the liquid "know" to evaporate when the volume was increased? The molecules in the liquid could not detectthe increase in volume for the gas, and thus could not possibly be responding to that increase.

The only reasonable conclusion is that the molecules in the liquid were always evaporating, even before thevolume of the container was increased. There must be a constant movement of molecules from the liquid phase into the gas phase.Since the pressure of the gas above the liquid remains constant when the volume is constant, then there must be a constant numberof molecules in the gas. If evaporation is constantly occurring, then condensation must also be occurring constantly, and moleculesin the gas must constantly be entering the liquid phase. Since the pressure remains constant in a fixed volume, then the number ofmolecules entering the gas from the liquid must be exactly offset by the number of molecules entering the liquid from the gas.