1. Changes of State
Chapter 10
Section 4

2. Review Summary- Physical States of Matter

3. Vocabulary
Phase: any part of a system that has uniform composition and properties
Condensation: the process by which a gas changes to a liquid
Equilibrium: a dynamic condition in which two opposing changes occur at equal rates in a closed system

4. More Information
Vapor: a gas in contact with its liquid or solid phase

6. Summary of State Changes See Table 2- p. 242

7. Equilibrium Vapor Pressure of a Liquid
Equilibrium vapor pressure: the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature
Increase temperature = increase in kinetic energy = increase in the number of molecules with enough energy to become a vapor = increase in equilibrium vapor pressure

8. At a Given Temperature
Because all liquids have characteristic forces of attraction between their particles, every liquid has a specific equilibrium vapor pressure at a given temperature
Stronger forces = smaller number of particles that can evaporate = low equilibrium vapor pressure

9. Relative Equilibrium Vapor Pressure
Volatile liquids: liquids that evaporate readily
Volatile liquids have weak forces of attraction between their particles.
Example: Ether, acetone, gasoline, methanol
Nonvolatile liquids: liquids that do not evaporate readily

10. Nonvolatile liquids have relatively strong attractive forces between their particles
Molten ionic compounds would be nonvolatile liquids because of the strong attraction for positive ions and negative ions in the liquid phase.

11. Boiling
Boiling allows the conversion of a liquid to a vapor within the liquid as well as at its surface.
Boiling point: the temperature at which the equilibrium vapor pressure of a liquid equals the atmospheric pressure
Energy must be added continuously for a liquid to keep boiling

12. Vapor Pressure Curves See Fig. 14- p. 344
normal

13. Pressure and Boiling Point
If you increase the pressure, then the boiling point is higher.
If you decrease the pressure, then the boiling point is lower.
At high altitudes, the atmospheric pressure is lower and water boils at a lower temperature, so it is recommended to bake at a higher temperature for less time. Some foods would require longer to cook.

15. The temperature at the boiling point remains constant despite the continuous addition of energy.
The added energy is used to overcome the attractive forces between the particles in the liquid as it vaporizes.
This energy is stored in the vapor as potential energy.

16. Molar Enthalpy of Vaporization
Molar enthalpy of vaporization: the amount of energy as heat required to evaporate 1 mole of a liquid at constant pressure and temperature
Abbreviated as
This is actually a measure of the strength of the attractive forces between particles
Each liquid has a characteristic molar enthalpy of vaporization

18. Freezing
Normal freezing point: the temperature at which the solid and liquid are in equilibrium at 1 atm (101.3 kPa)
Particles of the liquid and the solid have the same average kinetic energy
Energy lost during freezing is the loss of stored energy (potential)

19. Melting Melting is the reverse of freezing.
It ocurs at a constant temperature.
As a solid melts, it absorbs energy as heat.
For pure crystalline solids, the melting point and the freezing point are the same.
At equilibrium, melting and freezing proceed at equal rates.

20. Molar Enthalpy of Fusion
Molar enthalpy of fusion: the amount of energy as heat required to melt one mole of solid at the solid’s melting point
Abbreviated as
The energy absorbed overcomes the attractive forces and decreases particle order (dependent on the strength of the forces).

21. Sublimation and Deposition
Sublimation: the change of state from a solid directly to a gas
Examples: solid CO2 and iodine
Deposition: the change of state from a gas directly to a solid
Example: formation of frost on a cold surface

23. Figure 17-
p. 348
Blue Arrows = addition of energy to one phase to produce another
Green Arrows = removal of energy from one phase to produce another

24. Phase Diagrams
Phase diagram: a graph of pressure versus temperature that shows the conditions under which the phases of a substance exist
Triple point (of a substance): indicates the temperature and pressure conditions at which the solid, liquid, and vapor states can coexist at equilibrium

25. Critical point: of a substance, indicates the critical temperature and critical pressure
Critical temperature: the temperature above which the substance cannot exist in the liquid state
Critical pressure: the lowest pressure at which the substance can exist as a liquid at the critical temperature