This blog post will introduce the concepts of Thermodynamic phase changes of pure substances. Before proceeding, we need to define a pure substance as a substance with fixed chemical composition throughout. Note that a pure substance does not need to be a single chemical element. Water, for example, is a pure substance but contains more than one chemical element.
Phases and Thermodynamic Phase Changes
It is commonly known that substances can occur in different phases (solids, liquids, and gases). When substances are heated or cooled they can change from one phase to another. A phase change from a solid to a liquid is known as melting, and a phase change from a liquid to a gas is known as boiling (or vaporization). In the reverse direction, a phase change from a gas to a liquid is known as condensation and from a liquid to a solid is known as freezing.
The figure below illustrates phase changes in water plotted as temperature vs. heat addition. From state a to state b the water is a solid. Once the temperature reaches zero degrees Celsius (the melting point for water), adding more heat will result in a phase change. The phase change from solid to liquid is represented between points b and c on the diagram. The region between points c and d represent the liquid phase. More heat is added until the temperature reaches 100 degrees Celsius (the boiling point for water). As more heat is added the phase change will occur between a liquid and a gas, which is illustrated between points d and e. The region between points e and f is the gas phase. It is important to note that temperature remains constant during phase changes.
Consider now the phase diagram shown below, which is a plot of temperature (T) versus specific volume (v). The process from points 1 to 2 occur in a liquid phase, the phase change is shown between points 2 and 3, and from points 3 to 4 the substance is a gas. Notice again the transformation from one state to another, as shown in Figure 3 from state 2 to state 3, takes place with no change in temperature (horizontal line on the T-v diagram). The diagram will be used to define several important terms associated with phase change processes in pure substances.
When a substance occurs in a liquid state it is called a compressed liquid (or subcooled liquid). As heat is added to a liquid, it will reach a point where any additional heat added would result in some of the liquid vaporizing. A liquid at that point where it is about to vaporize is called a saturated liquid. State 2 represents the saturated liquid state on the figure.
The line between points 2 and 3 represent the substance in the phase change between a liquid and a vapor. That region is defined as a saturated liquid-vapor mixture, because both the liquid and vapor phase exists together in equilibrium.
A vapor that is not about to condense is known as a superheated vapor. As heat is removed from the vapor it will reach the point, known as a saturated vapor, where the removal of any additional heat will cause some of the vapor to condense. At a given pressure, the substance will begin to boil at a temperature known as the saturation temperature. Or, if temperature is held constant, the substance will start to boil at the saturation pressure.
The process line 1-2-3-4 occurs at a constant pressure. If the pressure is raised or lowered a similar process curve will exist, but the locations of the saturated liquid point and the saturated vapor point would change. The points for different pressures would form a curve defined by the saturated liquid line and the saturated vapor line.
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