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## Chapter 10: The Kinetic Theory of Matter

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Title: Chapter 10: The Kinetic Theory of Matter

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
• 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 273.15 K and boils at 373.15 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 0C and boils at 100C

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

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
ture-units-converting-between-kelvins-and-celsius.
htmllesson

8
• Examples
• 25C ? K 298 K ?C
• Tk (25 273) K Tc (298 - 273) C
• Tk 298 K Tc 25C

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
on-and-effusion-grahams-law.htmllesson

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

22

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 100C High altitude
96C
•
• 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