1. Thermodynamics
Phase Changes

2. Naming the Phase Changes
Phase changes occur because of changes in temperature and/or pressure.
Phase changes are always physical changes, as they do not involve any bonds or the creation of new substances.
Phase Change Names
Solid to liquid 
Melting
Liquid to solid
Freezing
Liquid to gas
Vaporization
Gas to liquid
Condensation
Solid to gas
Sublimation
Gas to solid
Deposition

3. Vapor Pressure
Some particles in a liquid or solid will have enough energy (KMT) to break away from the surface and become gaseous.
Vapor pressure is the pressure exerted by these molecules as they escape from the surface.
When the liquid or solid phase is in equilibrium with the gas phase, the pressure of the gas will be equal to the vapor pressure of the substance.
As temperature increases, the vapor pressure of the liquid will increase.
When the vapor pressure of a liquid increases to the point where it is equal to the surrounding atmospheric pressure, the liquid boils.

4. Heat of Fusion
The heat of fusion is the energy that must be put into a solid to melt it.
This is the energy needed to overcome the forces holding the solid together.
Alternatively, the heat of fusion (AKA heat of solidification) is the heat given off by a substance when it freezes.
Heat of fusion and solidification

5. Heat of Vaporization
The heat of vaporization is the energy that must be put into a liquid to turn it into a gas.
This energy is needed to overcome the forces holding the liquid together.
Alternatively, the heat of vaporization (AKA heat of condensation) is the heat given off by a substance when it condenses.
Heat of vaporization and condensation

6. Phase Diagrams
A phase diagram tells you the state of a substance at various pressure/temperature combinations.
In phase diagrams for substances other than water, the solid- liquid equilibrium line slopes upward.
Normally, when pressure is increased, a substance will change from liquid to solid.
In water, the line slopes downward, meaning water will change from solid to liquid.

7. Calorimetry
The specific heat of a substance is the amount of heat required to raise the temperature of one gram (or kg) of that substance by one degree Celsius (or Kelvin).
An object with a large specific heat can absorb a lot of heat without undergoing much of a temperature change, while the opposite is true of a substance with a low specific heat.
Calorimetry is the measurement of heat changes during chemical reactions, and it frequently is calculated using the following equation.
q = mcΔT q =
Heat added (J or cal) m
Mass of the substance (g or kg) c
Specific heat ΔT
Temperature change (K or oC)

8. Example
H+ + OH-  H2O(l) 25.0 mL of 1.5 M HCl and 30.0 mL of 2.0 M NaOH are mixed together in a Styrofoam cup and the reaction above occurs. The temperature of the reaction rises from 23.00oC to 31.60oC over the course of the reaction. Assuming the density of the solutions is 1.0 g/mL and the specific heat of the mixture is 4.18 J/goC, calculate the enthalpy of the reaction.

9. Heating Curves
A heating (or cooling) curve shows what happens to the temperature of a substance as heat is added.
If the substance is in a single phase, the temperature will increase.
The amount of the temperature increase can be calculated using calorimetry.
If the substance is in the process of a phase change, the temperature will remain constant.
The amount of heat to fully cause the substance to change phases can be calculated using the heat of fusion or the heat of vaporization.

10. Heating Curves

11. Example
A 1.53 g piece of ice is in a freezer and initially at a temperature of -15.1oC. The ice is removed from the freezer and melts completely after reaching a temperature of 0.0oC. If the specific heat of ice is 2.03 J/goC and its molar heat of fusion is 6.01 kJ/mol, how much heat is required for the entire process to occur?