1. States of Matter & Energy
Thermochemistry:
States of Matter & Energy

2. Phases of Matter From a solid to a liquid: Temperature increases
Attraction between particles decrease

3. Phases of Matter Solids Definite volume and shape
Attraction between molecules:
High
Ability for molecules to move:
Low (tend to vibrate in place)

4. Phases of Matter Liquid Definite volume; no definite shape
Attraction between molecules:
Medium
Ability for molecules to move:

5. Phases of Matter Gas No definite volume or shape
Attraction between molecules:
Low
Ability for molecules to move:
High

6. Phase Changes Sublimation Melting Boiling (fusion) (vaporization)
Freezing
Condensation
Deposition

7. Heating/Cooling Curve
Shows the temperature and energy of a substance over time as it changes from a solid to a gas
To increase the temperature we must add energy.

8. Energy
Temperature is a measure of the average kinetic energy of the particles.
Kinetic energy = energy of motion
Heat is the flow of energy from one object to another
As we add heat energy:
Kinetic energy of the particles increases, and temperature increases

9. Energy In solids, liquids, and gases:
As the average kinetic energy of the particles changes, the temperature changes
∆ KE = ∆ Temperature

10. Phase Changes Melting point Freezing point
Added heat energy allows particles to move around and overcome intermolecular attraction
As molecules move away from each other, solid becomes a liquid
There is no temperature change until all of the solid changes to a liquid
Freezing point
Heat is lost, and attractive forces pull molecules closer together to form solid
There is no temperature change until all of the liquid changes to a solid

11. Phase Changes
Heat energy during the melting phase is being used to overcome attractive forces between molecules
Solid phase: molecules stay in place and vibrate
Liquid phase: molecules can flow past each other
With a phase change, there is a change in potential energy (energy of position of molecules next to each other)

12. Phase Changes
Melting point and freezing point are at the same temperature
Melting: PE increases
Freezing: PE decreases
Boiling point and condensation point are the same temperature
Boiling: PE increases
Condensation: PE decreases
Each substance has its own characteristic melting/freezing point

13. Heat The amount of heat needed depends on:
Identify each phase and energy type (KE/PE) for sections a, c, and e: e d
Temp. c b a
Time
How much heat is needed to raise the temperature in a solid, liquid, or gas?
The amount of heat needed depends on:
How much substance you have (mass)
How much you want to change the temperature ( Temp)
The substance itself

14. Specific Heat q = mC T a: CpIce = 2.03 J/gC c: CpWater = 4.184 J/gCd
Temp. c b a
Time
q = mC T
q = heat
m = mass
T = change in temp.
C = specific heat
a: CpIce = 2.03 J/gC
c: CpWater = J/gC
e: CpWater Vapor = 2.00 J/gC

15. Specific Heat
How much heat is required to change 10.0 g of water from 20.0C to 50.0C?
q = m C T
q = (10.0 g) (4.184 J/gC)(50.0C C)
q = (10.0 g) (4.184 J/gC)(30.0C)
q = or 1260 J

16. Specific Heat
How much heat is required to change 10.0 g of ice from -30.0C to -10.0C?
q = m C T
q = (10.0 g) (2.03 J/gC)(-30.0C C)
q = (10.0 g) (2.03 J/gC)(20.0C)
q = 406 or 410 J

17. Heat of Fusion/Vaporization
Identify each phase and energy type (KE/PE) for sections b and d: e d
Temp. c b a
Time
How much heat is needed to change a solid to a liquid, or a liquid to a gas?
The amount of heat needed depends on:
How much substance you have (mass)
The substance itself

18. Heat of Fusion/Vaporizationd
Temp. c b a
Time
q = m H
q = heat
m = mass
H = heat of fusion or heat of vaporization
(No temp. change)
b: Hfus = 340 J/g
d: Hvap = 2260 J/g

19. Heat of Fusion How much heat is needed to melt 25.0 g of ice?
q = m Hfus
q = (25.0 g)(340 J/g)
q = 8500 J

20. Heat of Vaporization
How much heat is needed to evaporate 25.0 g of water?
q = m Hvap
q = (25.0 g)(2260 J/g)
q = J

21. Heat of Solution Heat of solution
The heat produced by a chemical reaction, or the heat required for a chemical reaction to occur
Heat is absorbed from the atmosphere or released into the atmosphere

22. Heat of Solution Endothermic reaction Exothermic reaction
Requires heat energy from the environment to get reaction to run
Heat is transferred from the environment to the reaction
Is a positive value
Exothermic reaction
Produces heat
Heat is transferred from the reaction to the environment
Is a negative number

23. Heat of Solution problem
What is the enthalpy of solution in KJ/mol if 6.34 g of KOH is dissolved into 98.4 g of water. The temperature of the water changes from 24.5oC to 22.0oC. Is this exothermic or endothermic?

24. Boiling Point Does water always boil at 100ºC? Boiling point: NO.
Boiling point depends on atmospheric pressure
Boiling point:
Is the temperature at which vapor pressure of the liquid is the same as the atmospheric pressure
Pressure of the bubbles = air pressure in the room