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By the end of this section, you will be able to:
  • Describe phase transitions and equilibrium between phases
  • Solve problems involving latent heat
  • Solve calorimetry problems involving phase changes

Phase transitions play an important theoretical and practical role in the study of heat flow. In melting (or “fusion”), a solid turns into a liquid; the opposite process is freezing . In evaporation , a liquid turns into a gas; the opposite process is condensation .

A substance melts or freezes at a temperature called its melting point, and boils (evaporates rapidly) or condenses at its boiling point. These temperatures depend on pressure. High pressure favors the denser form, so typically, high pressure raises the melting point and boiling point, and low pressure lowers them. For example, the boiling point of water is 100 ° C at 1.00 atm. At higher pressure, the boiling point is higher, and at lower pressure, it is lower. The main exception is the melting and freezing of water, discussed in the next section.

Phase diagrams

The phase of a given substance depends on the pressure and temperature. Thus, plots of pressure versus temperature showing the phase in each region provide considerable insight into thermal properties of substances. Such a pT graph is called a phase diagram    .

[link] shows the phase diagram for water. Using the graph, if you know the pressure and temperature, you can determine the phase of water. The solid curves—boundaries between phases—indicate phase transitions, that is, temperatures and pressures at which the phases coexist. For example, the boiling point of water is 100 ° C at 1.00 atm. As the pressure increases, the boiling temperature rises gradually to 374 ° C at a pressure of 218 atm. A pressure cooker (or even a covered pot) cooks food faster than an open pot, because the water can exist as a liquid at temperatures greater than 100 ° C without all boiling away. (As we’ll see in the next section, liquid water conducts heat better than steam or hot air.) The boiling point curve ends at a certain point called the critical point    —that is, a critical temperature    , above which the liquid and gas phases cannot be distinguished; the substance is called a supercritical fluid . At sufficiently high pressure above the critical point, the gas has the density of a liquid but does not condense. Carbon dioxide, for example, is supercritical at all temperatures above 31.0 ° C . Critical pressure is the pressure of the critical point.

Graph of pressure P in atmosphere versus temperature T in degree Celsius for water. The curve starts with going up and right to a point labeled triple point. This is at 0.006 atm and 0.01 degree C. From here, the curve diverges into two branches. One goes up and left and is almost vertical. The other goes up and right. On the branch going up and right is a point at 1 atm and 100 degrees C. Further up on the same branch is a point labeled critical point. This is at 218 atm and 374 degrees C. The area to the left of the left branch is labeled solid. The area between two branches is labeled liquid. The area to the right of the right branch is labeled vapour. The curve to the lower left of the triple point is labeled sublimation, the branch to the upper left of the triple point is labeled melting, and the branch to the upper right of the triple point is labeled boiling.
The phase diagram ( pT graph) for water shows solid (s), liquid (l), and vapor (v) phases. At temperatures and pressure above those of the critical point, there is no distinction between liquid and vapor. Note that the axes are nonlinear and the graph is not to scale. This graph is simplified—it omits several exotic phases of ice at higher pressures. The phase diagram of water is unusual because the melting-point curve has a negative slope, showing that you can melt ice by increasing the pressure.

Similarly, the curve between the solid and liquid regions in [link] gives the melting temperature at various pressures. For example, the melting point is 0 ° C at 1.00 atm, as expected. Water has the unusual property that ice is less dense than liquid water at the melting point, so at a fixed temperature, you can change the phase from solid (ice) to liquid (water) by increasing the pressure. That is, the melting temperature of ice falls with increased pressure, as the phase diagram shows. For example, when a car is driven over snow, the increased pressure from the tires melts the snowflakes; afterwards, the water refreezes and forms an ice layer.

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Source:  OpenStax, University physics volume 2. OpenStax CNX. Oct 06, 2016 Download for free at http://cnx.org/content/col12074/1.3
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