0.12 Phase equilibrium and intermolecular interactions  (Page 5/7)

Recall that, at the boiling point, we observe that both liquid and gas are at equilibrium with one another. Thisis true at every combination of applied pressure and boiling point temperature. Therefore, for every combination of temperature andpressure on the graph in , we observe liquid-gas equilibrium.

What happens at temperature/pressure combinations which are not on the line in ? To find out, we first start at a temperature-pressure combination on the graph and elevate thetemperature. The vapor pressure of the liquid rises, and if the applied pressure does not also increase, then the vapor pressurewill be greater than the applied pressure. We must therefore not be at equilibrium anymore. All of the liquid vaporizes, and there isonly gas in the container. Conversely, if we start at a point on the graph and lower the temperature, the vapor pressure is belowthe applied pressure, and we observe that all of the gas condenses into the liquid.

Now, what if we start at a temperature-pressure combination on the graph and elevate theapplied pressure without raising the temperature? The applied pressure will be greater than the vapor pressure, and all of thegas will condense into the liquid. thus actually reveals to us what phase or phases are present at each combination of temperature andpressure: along the line, liquid and gas are in equilibrium; above the line, only liquid is present; below the line, only gas ispresent. When we label the graph with the phase or phases present in each region as in , we refer to the graph as a phase diagram .

Of course, only includes liquid, gas, and liquid-gas equilibrium. We know that, if the temperature is lowenough, we expect that the water will freeze into solid. To complete the phase diagram, we need additional observations.

We go back to our apparatus in and we establish liquid-gas water phase equilibrium at a temperature of 25°C and 23.8 torr. Ifwe slowly lower the temperature, the vapor pressure decreases slowly as well, as shown in . If we continue to lower the temperature, though, we observe an interesting transition, as shownin the more detailed . The very smooth variation in the vapor pressure shows a slight, almostunnoticeable break very near to 0°C. Below this temperature, the pressure continues to vary smoothly, but along a slightlydifferent curve.

To understand what we have observed, we examine the contents of the container. We find that, attemperatures below 0°C, the water in the container is now an equilibrium mixture of water vapor and solid water (ice), and thereis no liquid present. The direct transition from solid to gas, without liquid, is called sublimation . For pressure-temperature combinations along this new curve below0°C, then, the curve shows the solid-gas equilibrium conditions. As before, we can interpret this two ways. Thesolid-gas curve gives the vapor pressure of the solid water as a function of temperature, and also gives the sublimation temperatureas a function of applied pressure.