The document summarizes key concepts about phase changes and energy changes, including: 1) Phase changes can be exothermic or endothermic. 2) Heating and cooling curves illustrate changes in kinetic and potential energy that occur during temperature and phase changes. 3) The specific heat capacity of a substance determines how much its temperature changes with added or removed energy.

1. Unit 7b Honors Chemistry Review: Phase Changes and Energy Changes
You should be able to:
1. Identify exchanges of energy as exothermic or endothermic during phase changes. (pg 8)
2. Interpret a triple point diagram.
Special note: A point that falls on any of the lines is where the two phases are in equilibrium. This means that
you have both phases present.
3. Define critical point and triple point.
Critical point is the temperature and pressure at which the properties of the vapor phase of a substance can’t
be distinguished from those of the liquid phase.
Triple point is the temperature and pressure at which the solid, liquid, and gas phases of a substance exist in
equilibrium (all are present at the same time).
4. Tell how temperature relates to kinetic energy. (pg 44)
Temperature is a measure of the average kinetic energy of a substance. As temperature increases, kinetic
energy increases. As temperature decreases, kinetic energy decreases. This is a direct relationship.
5. Interpret a heating and cooling curve. (notes)
Heating Curve Cooling Curve

2. 6. List the names of the phase changes matter can undergo. (pgs 454-458, and notes)
Vaporization: liquid to gas Condensation: gas to liquid
Melting: solid to liquid Freezing: liquid to solid
Sublimation: solid to gas Deposition: gas to solid
7. Describe changes in kinetic and potential energy during changes in temperature and phase using heating and
cooling curves. (notes)
When temperature is changing kinetic energy (KE) is changing. If temperature is increasing, KE is
increasing. If temperature is decreasing, KE is decreasing. If temperature is not changing, KE is not
changing.
When temperature is constant, potential energy (PE) is changing. If energy is being added but the
temperature is not increasing, then PE is increasing. If energy is being removed as in a cooling curve, but
temperature is not changing, PE is decreasing.
8. Define specific heat capacity (Cp).
Specific heat capacity is the amount of heat energy required to raise the temperature of 1.0 gram of a
substances by 1 Kelvin (or 1 degree Celsius)
9. Describe how specific heat and changes in temperature are related. (pg 45)
If you add the same amount of heat energy to similar masses of different substances, they don’t all show the
same increase in temperature. Think about a park bench made out of metal and a park bench made out of wood.
If they are sitting side by side in the sun, which bench will be hotter?
The metal bench will be hotter, because it has a lower specific heat capacity. This means that it takes less
energy to raise the temperature. The lower the specific heat capacity, the quicker a substance heats up!
Iron 0.449 J/g-K Oak 2.00 J/g-k
10. Calculate energy change during a temperature change using Q= mCp∆T. (notes)
Q = heat energy
m= mass (normally in grams, but check units on Cp)
∆T = change in temperature (Tfinal - Tinital)
11. Calculate energy changes during a phase changes using E=mol∆Hfus or E = mol∆Hvap (notes)
E = heat energy
Mol = # of moles (if given grams, you must convert grams to moles by dividing by the molar mass)
∆Hfus = heat of fusion (energy needed for the phase change of solid to liquid or energy released during the
phase change liquid to solid)
∆Hvap = heat of vaporization (energy needed for the phase change of liquid to a gas or energy released
during the phase change gas to a liquid)

3. 12. Describe the relationship between elevation and atmospheric pressure (pg 456) and how changes in atmospheric
pressure affect the boiling point. (pg 45)
As you go up in elevation, atmospheric pressure increases/decreases and the boiling point of a substance
increases, decreases.
Water boils at a higher/lower temperature at sea level (1 atm) than in the Rocky Mountains (0.90 atm)
13. Define normal boiling point and vapor pressure. (pgs 454-456)
Normal boiling point is the temperature at which a substance boils at 1.00 atm of pressure.
Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid state at a given temperature.
(Note: As temperature increases, vapor pressure increases.)
14. When does a liquid begin to boil? (pg 456)
A liquid begins to boil when the external pressure on the liquid equals the vapor pressure of the liquid. If
the container is open, the external pressure is typically atmospheric pressure.
Therefore if you lower the atmospheric pressure, the boiling point of the liquid decreases.
15. Compare evaporation to boiling. (pgs 399-400)
Evaporation is when a liquid turns into a gas at a temperature lower than the boiling point of the liquid and
occurs on the surface of the liquid.
Boiling is when a liquid turns into a gas at the boiling point of the liquid and originates within the liquid.
Practice Problems:
1. If the specific heat of water is 4.18 J/g°C, how much energy would be needed to heat 20 grams of water from
the freezing point to the boiling pt? Is this an exothermic or endothermic change?
2. If the heat of fusion for ice is 6.0 kJ/mol, how many moles of ice can be melted with 325.3 J of energy?
3. What happens to the motion of the particles in a gas as temperature is increased? What type of energy is
changing when this happens?
4. What happens to the particles in a substance during a phase change? What type of energy is changing when
this happens?