PHASE CHANGES
A. Phase- part of a system that has uniform properties and composition; phase changes are physical
changes based on heat content.
B. Common Phase Changes
1. melting- solid to liquid
2. sublimation- solid to gas
3. freezing- liquid to solid
4. vaporization- liquid to a gas
a. evaporation- liquid particles escape the surface of a nonboiling liquid as a gas.
b. boiling- conversion of liquid to gas within the liquid and at the surface when the
vapor pressure equals the atmospheric pressure; attractive forces are overcome.
5. condensation- gas to liquid
6. deposition- gas to solid
C. Additional Phase Concepts
1. supercooled liquids- substances that retain certain liquid properties even at temperatures at
which they appear to be solid.
2. gas liquefication- Michael Faraday 1823, English chemist; cooling and compressing gases at
the same time would yield a liquid; a critical point for temperature and pressure. Faraday
liquefied chlorine, carbon dioxide, hydrogen sulfide, and hydrogen bromide.
3. volatile liquids- liquids that easily evaporate due to weak intermolecular forces.
4. molar heat of vaporization- amount of heat energy needed to vaporize one mole of liquid at its
boiling point.
5. molar heat of fusion- amount of heat energy needed to melt one mole of solid at its melting
point
D. Equilibrium- point where two opposing changes occur at equal rates in a closed system.
1. closed system- no substances are added or lost, but energy changes freely occur.
2. physical equilibrium- a state in which two opposing physical changes occur at equal rates in the
same system.
3. equilibrium vapor pressure- the pressure exerted by a vapor in equilibrium with its liquid.
4. dynamic equilibrium- a forward and back reaction reach a point where there is no further
change in the quantities of substances, but the reactions are still taking place. This is shown by
a reversible reaction arrow. (
)
5. Le Chatelier’s Principle- Henri Le Chatelier 1884
a. If a system is subjected to stress, the equilibrium will change in order to relieve the
stress.
b. equation will shift based on concentration, partial pressure of gas (number of
molecules is important), and temperature; consider if the reaction is endothermic
(H is positive) or exothermic (H is negative); (see figure 1)
c. catalysts do not change equilibrium, they only speed up the rate at which the reaction
reaches dynamic equilibrium
E. Phase Relationships
1. phase diagram- a graph of pressure vs. temperature that shows the conditions under which the
phases of a substance exist. (see figures 2 and 3)
a. triple point- the temperature and pressure conditions at which the solid, liquid, and
vapor of the substance can coexist at equilibrium
b. critical point- indicates a critical temperature and critical pressure
c. critical temperature- the temperature above which the substance cannot exist in the
liquid state
d. critical pressure- the lowest pressure at which the substance can exist as a liquid at
the critical temperature
2. supercritical fluid- a material that can be either liquid or gas, used above the critical point
where gases and liquids can coexist. This fluid has gaseous properties of penetrating
substances, and liquid properties of dissolving materials into their components.
Figure 1: Le Chatelier’s Equilibrium
Concentration
Pressure
Temperature (Exothermic Reaction)
Figure 2: General Phase Diagram
Figure 3: Specific Phase Diagrams