1. Chapter 10: The Kinetic Theory of Matter
Section 10.2: Kinetic Energy and Changes of State

2. Main Idea: Matter changes states when energy is added or removed
A) Interpret changes in temperature and changes of state of a substance in terms of the kinetic theory of matter
B) Relate Kelvin and Celsius Temperature scales
C) Analyze the effects of temperature and pressure on changes of state

3. Temperature and Kinetic Energy
Particles move in random directions at different rates
Temperature- measure of the average kinetic energy of the particles that make up the material
As gas is heated, the average kinetic energy and speed of its particles increases → temperature increases
As gas is cooled, the average kinetic energy and speed of its particles decreases → temperature decreases

4. Kelvin Scale
Kelvins= SI unit of temperature (divisions on a Kelvin scale)
Water freezes at K and boils at K
The temperature at which a substance would have zero kinetic energy is called Absolute zero
Absolute zero has never been reached because submicroscopic particles are in constant motion
It is defined so that the temperature of a substance is directly proportional to the average kinetic energy of the particles and so the zero on the Kelvin scale corresponds to zero kinetic energy

5. Celsius Scale used throughout the world
Water freezes at 0°C and boils at 100°C

6. Fahrenheit Scale used by weather reporters, household ovens
Water freezes at 32°F and boils at 212°F

7. Temperature Conversions
The divisions of the Fahrenheit and Celsius are called degrees, but the divisions of the Kelvin scale are called Kelvins
Celsius to Kelvin:
Tk = (Tc +273) K
Kelvin to Celsius:
Tc = (Tk - 273) °C

8. Examples:
25°C = ? K K =?°C
Tk = ( ) K Tc = ( ) °C
Tk = 298 K Tc = 25°C

9. Changes of State Dependent on temperature
Include: Evaporation, Sublimation, Condensation, Melting, Freezing, Deposition

10. Evaporation
the process by which particles of a liquid form a gas by escaping from the surface
The area of the surface, temperature, and humidity affect the rate of evaporation
Liquids that evaporate quickly are volatile liquids
Example: perfume

11. Sublimation process by which solid goes to a gas
occurs when the solid to liquid state is skipped
Example of material that sublimes: dry ice (solid CO2)

12. Deposition
The opposite of sublimation
Gas goes into a solid

13. Condensation the reverse of evaporation (gas → liquid)
the gas particles come closer together (condense) and form a liquid

14. Melting Point
the temperature of the solid when its crystal lattice begins to disintegrate
When more heat is applied after the melting point, energy is used until the crystal lattice collapses and becomes a liquid

15. Freezing Point
If a liquid substance is cooled, the temperature falls, and the liquid becomes a solid
The temperature of a liquid when it begins to form a crystal lattice and becomes a solid

16. During Phase Changes…
Because energy is always conserved, energy is released when vapor changes to a liquid
As with boiling and condensing, the kinetic energies of the particles of a substance do not change during melting or freezing

17. Mass and Speed of Particles
Particles of greater mass have greater kinetic energy
Particles with greater speed have greater kinetic energy
Motions of gas particles cause them to spread out to fill containers uniformly

18. Diffusion
process by which particles of matter fill a space because of random motion (ex: food coloring moving in water)
The rate of diffusion of a gas depends upon its kinetic energy- mass and speed of its molecules

19. Vapor Pressure
The liquid water that is left in a closed container will not all evaporate. The liquid in a closed container comes to equilibrium with its vapor
When equilibrium is reached, the pressure exerted by vapors reaches its final, maximum value (volume of liquid will not change)

20. Vapor Pressure (cont)
Vapor pressure - The pressure of a substance in equilibrium with its liquid (rates of evaporation and condensation are equal)
The value of vapor pressure of a substance indicates how easily the substance evaporates:
High vapor pressure = more volatile
Low vapor pressure = less volatile
Higher temperatures = greater vapor pressure
Lower temperatures = less vapor pressure

21. Boiling point
Temperature of the substance when its vapor pressure equals the pressure exerted in on the surface of the liquid
Normal boiling point is the temperature at which liquid boils in an open container at normal atmospheric pressure

23. Boiling point (cont)
Boiling point of a liquid increases when pressure increases
Boiling point of a liquid decreases when pressure decreases
Example: Sea level: 100°C High altitude: 96°C
Because the temp of the boiling water is lower at high elevations, it takes longer to cook foods.
↑ altitude = ↓ pressure = ↓ boiling point

24. Heat of Vaporization
energy absorbed when 1 kg of a liquid vaporizes at its normal boiling point
Joule (J) -SI unit of energy required to lift a 1-kg mass 1 meter against the force of gravity
2.26 x 106 J is the energy needed to move molecules in 1 kg of water far enough apart that they form water vapor
Heat of vaporization of water = 2.26 x 106 J/kg

25. Heat of Vaporization (cont)
Example: How much energy is absorbed if a 500g sample of water vaporizes?
0.5 Kg x 2.26 x 106J = 1.13 x 106 J
1 Kg

26. Heat of Fusion
The energy released as 1 kg of a substance solidifies at its freezing point
Heat of fusion of water = 3.34 x 105 J/kg

27. Heat of Fusion (cont)
Example: How much energy is released if a 5000g sample of water solidifies?
5 Kg x 3.34 x 105J = 1.67 x 106 J
1 Kg

28. Heating and Cooling curves