However, the solid-liquid phase boundary for water is anomalous, in that it has a negative slope. This reflects the fact that ice is lower in density than liquid water a well-known fact: ice floats , unlike most other substances which typically have denser solid phases.
Mothballs — Phase Transitions Applied : The thermodynamic properties of mothballs, made of 1,4-Dichlorobenzene, are used to repel insects.
The gas released is toxic to insects. Privacy Policy. Skip to main content. Liquids and Solids. Search for:. Phase Diagrams Major Features of a Phase Diagram Phase diagrams are useful because they allow us to understand in what state matter exists under certain conditions. Learning Objectives Describe the major features of a phase diagram. Key Takeaways Key Points The major features of a phase diagram are phase boundaries and the triple point.
Phase diagrams demonstrate the effects of changes in pressure and temperature on the state of matter. At phase boundaries, two phases of matter coexist which two depends on the phase transition taking place. The triple point is the point on the phase diagram at which three distinct phases of matter coexist in equilibrium.
Key Terms Triple point : The unique temperature and pressure at which the solid, liquid, and gas phases of a substance are all in equilibrium with each other. Interpreting Phase Diagrams Phase diagrams illustrate the effects selected variables of a system have on the state of matter. We call the gas phase a vapor when it exists, as it does for water at The ice and liquid water are in thermal equilibrium, so that the temperature stays at the freezing temperature as long as ice remains in the liquid.
Once all of the ice melts, the water temperature will start to rise. Vapor pressure is defined as the pressure at which a gas coexists with its solid or liquid phase. Vapor pressure is created by faster molecules that break away from the liquid or solid and enter the gas phase.
The vapor pressure of a substance depends on both the substance and its temperature—an increase in temperature increases the vapor pressure. Partial pressure is defined as the pressure a gas would create if it occupied the total volume available. In a mixture of gases, the total pressure is the sum of partial pressures of the component gases , assuming ideal gas behavior and no chemical reactions between the components.
Thus water evaporates and ice sublimates when their vapor pressures exceed the partial pressure of water vapor in the surrounding mixture of gases. If their vapor pressures are less than the partial pressure of water vapor in the surrounding gas, liquid droplets or ice crystals frost form.
Is energy transfer involved in a phase change? If so, will energy have to be supplied to change phase from solid to liquid and liquid to gas? What about gas to liquid and liquid to solid? Why do they spray the orange trees with water in Florida when the temperatures are near or just below freezing? Yes, energy transfer is involved in a phase change. We know that atoms and molecules in solids and liquids are bound to each other because we know that force is required to separate them.
So in a phase change from solid to liquid and liquid to gas, a force must be exerted, perhaps by collision, to separate atoms and molecules. Force exerted through a distance is work, and energy is needed to do work to go from solid to liquid and liquid to gas. This is intuitively consistent with the need for energy to melt ice or boil water. The converse is also true. Going from gas to liquid or liquid to solid involves atoms and molecules pushing together, doing work and releasing energy.
Heat, cool, and compress atoms and molecules and watch as they change between solid, liquid, and gas phases. Figure 5. The phase diagram for carbon dioxide. The axes are nonlinear, and the graph is not to scale. Dry ice is solid carbon dioxide and has a sublimation temperature of — Skip to main content. Temperature, Kinetic Theory, and the Gas Laws. Search for:.
Phase Changes Learning Objectives By the end of this section, you will be able to: Interpret a phase diagram. Identify and describe the triple point of a gas from its phase diagram. Describe the state of equilibrium between a liquid and a gas, a liquid and a solid, and a gas and a solid.
Solution The ice and liquid water are in thermal equilibrium, so that the temperature stays at the freezing temperature as long as ice remains in the liquid. Check Your Understanding Is energy transfer involved in a phase change?
Solution Yes, energy transfer is involved in a phase change. PhET Explorations: States of Matter—Basics Heat, cool, and compress atoms and molecules and watch as they change between solid, liquid, and gas phases.
Click to download the simulation. Run using Java. Very careful measurements reveal that the solid-gas line and the liquid-gas line intersect in Figure Under these conditions, we observe inside the container that solid, liquid, and gas are all three at equilibrium inside the container.
As such, this unique temperature-pressure combination 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 that the temperature at which the solid and liquid are in equilibrium varies almost imperceptibly as we increase the pressure. If we include the solid-liquid equilibrium conditions on the previous phase diagram, we get Figure Each substance has its own unique phase diagram, corresponding to the diagram in Figure There are several questions raised by our observations of phase equilibrium and vapor pressure.
The first we will consider is why the pressure of a vapor in equilibrium with its liquid does not depend on the volume of the container into which the liquid evaporates, or on the amount of liquid in the container, or on the amount of vapor in the container. Why do we get the same pressure for the same temperature, regardless of other conditions?
To address this question, we need to understand the coexistence of vapor and liquid in equilibrium. How is this equilibrium achieved? To approach these questions, let us look again at the situation in Figure 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 larger fixed 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 have evaporated 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 molecules move.
How did the liquid "know" to evaporate when the volume was increased? The molecules in the liquid could not detect the 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 the volume 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 number of molecules in the gas.
If evaporation is constantly occurring, then condensation must also be occurring constantly, and molecules in the gas must constantly be entering the liquid phase. Since the pressure remains constant in a fixed volume, then the number of molecules entering the gas from the liquid must be exactly offset by the number of molecules entering the liquid from the gas.
At equilibrium, therefore, the pressure and temperature inside the container are unchanging, but there is constant movement of molecules between the phases. This is called dynamic equilibrium.
The situation is "equilibrium" in that the observable properties of the liquid and gas in the container are not changing, but the situation is "dynamic" in that there is constant movement of molecules between phases. The dynamic processes that take place offset each other exactly so that the properties of the liquid and gas do not change. What happens when we increase the volume of the container to a larger fixed volume?
We know that the pressure equilibrates at the same vapor pressure, and that therefore there are more molecules in the vapor phase. How did they get there? It must be the case that when the volume is increased, evaporation initially occurs more rapidly than condensation until equilibrium is achieved.
The rate of evaporation must be determined by the number of molecules in the liquid which have sufficient kinetic energy to escape the intermolecular forces in the liquid, and according to the kinetic molecular theory, this number depends only on the temperature, not on the volume.
However, the rate of condensation must depend on the frequency of molecules striking the surface of the liquid. According to the Kinetic Molecular Theory, this frequency must decrease when the volume is increased, because the density of molecules in the gas decreases. Therefore, the rate of condensation becomes smaller than the rate of evaporation when the volume is increased, and therefore there is a net flow of molecules from liquid to gas.
This continues until the density of molecules in the gas is restored to its original value, at which point the rate of evaporation is matched by the rate of condensation. At this point, this pressure stops increasing and is the same as it was before the volume was increased. In the phase diagram for water in Figure State what happens physically to the water during this heating process. State what happens physically to the water during this process.
Explain why Figure We observe that, when the applied pressure is less than the vapor pressure of a liquid, all of the liquid will spontaneously evaporate. In terms of dynamic equilibrium, explain why no liquid can be present under these conditions. Using arguments from the Kinetic Molecular Theory and the concept of dynamic equilibrium, explain why, at a given applied pressure, there can be one and only one temperature, the boiling point, at which a specific liquid and its vapor pressure can be in equilibrium.
Using dynamic equilibrium arguments, explain why the vapor pressure of a liquid is independent of the amount of liquid present. Using dynamic equilibrium arguments, explain why the vapor pressure of a liquid is independent of the volume available for the vapor above the liquid. Using dynamic equilibrium arguments, explain why a substance with weaker intermolecular forces has a greater vapor pressure than one with stronger intermolecular forces.
According to Figure Which of these substances has the greater intermolecular attractions? Which substance has the higher boiling point? Explain the difference in the intermolecular attractions in terms of molecular structure. The text describes dynamic equilibrium between a liquid and its vapor at the boiling point.
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