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Phases of Matter, Phase Transitions, Latent vs Sensible Heat, Phase Diagrams

Phases of matter refer to the distinct forms in which a substance can exist, such as solid, liquid, or gas. These phases are characterized by their physical states and the behavior of their particles. Phase transitions are the processes by which a substance changes from one phase to another, such as melting, freezing, vaporization, or condensation. These transitions occur when the conditions, like temperature or pressure, change enough to alter the arrangement and motion of the particles. Understanding phases and phase transitions is essential in thermodynamics because it helps explain how energy is stored and transferred during physical changes. This topic also introduces the concepts of latent heat and sensible heat, which describe the energy involved in phase changes versus temperature changes, and phase diagrams, which visually represent the conditions under which different phases exist. These ideas are fundamental to understanding the behavior of matter in various scientific and engineering contexts.

Figures (7)

Figure 1.12 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.
Figure 1.13 Direct transitions between solid and vapor are common, sometimes useful, and even beautiful. (a) Dry ice sublimes directly to carbon dioxide gas. The visible “smoke” consists of water droplets that condensed in the air cooled by the dry ice. (b) Frost forms patterns on a very cold window, an example of a solid formed directly from a vapor. (credit a: modification of work by Windell Oskay; credit b: modification of work by Liz West)
Figure 1.14 Equilibrium between liquid and gas at two different boiling points inside a closed container. (a) The rates of boiling and condensation are equal at this combination of temperature and pressure, so the liquid and gas phases are in equilibrium. (b) At a higher temperature, the boiling rate is faster, that is, the rate at which molecules leave the liquid and enter the gas is faster. This increases the number of molecules in the gas, which increases the gas pressure, which in turn increases the rate at which gas molecules condense and enter the liquid. The pressure stops increasing when it reaches the point where the boiling rate and the condensation rate are equal. The gas and liquid are in equilibrium again at this higher temperature and pressure.
Figure 1.15 Temperature versus heat. The system is constructed so that no vapor evaporates while ice warms to become liquid water, and so that, when vaporization occurs, the vapor remains in the system. The long stretches of constant temperatures at 0°C0°C and 100°C100°C reflect the large amounts of heat needed to cause melting and vaporization, respectively.
Figure 1.16 Condensation forms on this glass of iced tea because the temperature of the nearby air is reduced. The air cannot hold as much water as it did at room temperature, so water condenses. Energy is released when the water condenses, speeding the melting of the ice in the glass. (credit: Jenny Downing)
Figure 1.17 The ice on these trees released large amounts of energy when it froze, helping to prevent the temperature of the trees from dropping below 0°C0°C. Water is intentionally sprayed on orchards to help prevent hard frosts. (credit: Hermann Hammer)
Figure 1.18 (a) Energy is required to partially overcome the attractive forces (modeled as springs) between molecules in a solid to form a liquid. That same energy must be removed from the liquid for freezing to take place. (b) Molecules become separated by large distances when going from liquid to vapor, requiring significant energy to completely overcome molecular attraction. The same energy must be removed from the vapor for condensation to take place.

Latent heat is the energy required to change a substance from one phase to another without changing its temperature. For example, when water boils, it absorbs latent heat to turn into steam, but its temperature stays at 100°C until all the water has vaporized. Sensible heat, on the other hand, is the energy that changes the temperature of a substance without changing its phase.

If you heat water from 20°C to 80°C, the energy you add is sensible heat. The key difference is that latent heat is used to break or form bonds between particles during a phase change, while sensible heat increases the motion of particles, raising the temperature. A student should picture latent heat as the hidden energy that causes a phase change, and sensible heat as the energy that makes a substance feel hotter or colder.

The diagram has three main regions: solid, liquid, and gas. Lines between these regions represent phase transitions, like melting or boiling. The point where all three regions meet is called the triple point, where all three phases coexist.

The critical point marks the end of the liquid-gas boundary, beyond which the substance becomes a supercritical fluid. A student should picture a phase diagram as a map showing how a substance behaves under different conditions. By following the lines, you can predict what phase a substance will be in at any given temperature and pressure.

To understand these ideas, it helps to think about water. At room temperature and normal pressure, water is a liquid. If you lower the temperature, it freezes into a solid.

Key Points

  • A solid is a phase that resists deformation and maintains a fixed shape under applied shear stress.
  • A liquid is a fluid phase with a relatively high density that is insensitive to changes in temperature and pressure.
  • A gas is a fluid phase with a relatively low density that is sensitive to changes in temperature and pressure.
  • Melting is an equilibrium phase transition from solid to liquid at a specific temperature and pressure.
  • Freezing is an equilibrium phase transition from liquid to solid at a specific temperature and pressure.

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