1. Prepare for Test TURN CELL PHONES OFF AND PLACE IN BACKPACK
Put your backpack at the front of the classroom
YOUR WILL NEED:
2 sheets of paper
Pencil
Calculator

2. Chapter 12- States of Matter

3. Chapter 12 - Essential Questions
Copy down these questions:
Unit EQ
What makes solids, liquids, and gases different from each other?
Lesson EQ’s
What is vapor pressure and how does that relate to boiling point?
What makes water a liquid and carbon dioxide a gas?
What is the Kinetic-Molecular Theory?

4. Chapter 12 Word Study Guide DUE Monday
Word, definition, your definition, picture
Gas pressure
Kinetic Molecular theory
Elastic Collision
Dalton’s Law of Partial Pressure
Condensation
Critical point
Deposition
Freezing
Melting
Melting point
Phase change diagram of water (temp vs. heat)
Phase diagram (pressure vs. temp)
Sublimation
Triple point
Vaporization
Covalent Network
Dipole to dipole forces
Hydrogen bonding
Ionic
London Dispersion Forces

5. Journal Review the following Concepts
Write down everything you know about the following topics (Include information from your textbook)
Density
Physical Properties
Three common states of matter
Conversions
Kelvin  Celsius
Celsius  Farenheit

6. Convert the following temperatures
How many degrees Kelvin is 45°C?
How many degrees Celsius is 303 K
Convert 75°F to Celsius
Convert 40°C to Fahrenheit

7. Page 400 Take 2 or 3 minutes to analyze this picture Questions
What is it describing?
Questions
What are the circles in the photo?
Why are they different shades of purple?
What makes them different shades of purple?
What is the actual container being used?

8. Gases What determines an atoms Chemical Properties?
Are physical properties also affected?

9. Kinetic Molecular Theory
Ludwig Boltzmann and James Maxwell
Worked independently
Developed a model to explain properties of gases
KINETIC MOLECULAR THEORY
Describes the behavior of matter in terms of particles in motion.
Size
Motion
Energy

10. Particle Size Small particles
So far a part the they experience almost NO attractive or repulsive forces.
What brings the atoms together?

11. Particle Motion Particles are in constant random motion.
WHY?????
Move in a straight line until they collide with another particle or a wall
Elastic = NO kinetic energy is lost
May be transferred but the total energy is the same.

12. Particle Energy KE = ½ mv2
V= velocity (speed and direction)
M = mass
All particles have the same mass but not same velocity ( )
Temperature
Measure of the average kinetic energy of the particles in a sample of matter.

13. Journal
How does temperature affect the speed at which and object would sink in it?
Example: a marble being placed in oil.

14. Explaining the Behavior of Gases
Gas particles expand until the container is filled
Three main explanations
Low density
Compression and Expansion
Diffusion and Effusion

15. Low Density D = m/v In your book
Gold is 6500 times more dense than Chlorine
What does this mean???

16. Compression and Expansion
Compress = reduce volume
Expand = increase volume
Why can air be compressed?
How does density relate to change in pressure?
Look at figure 12.3
Individually write down your thoughts.

17. Diffusion and Effusion
The movement of particles from a high concentration to a low concentration
What kinds of particles are in the air?
If I sprayed something, what would happen?
How can they move through air that is already filled with another particle?
Effusion
When gas escapes through a tiny opening.

18. Diffusion

19. Effusion

20. Rate of Effusion Discovered by Thomas Graham
Rate of Effusion = 1/√molar mass
Example
What is the rate of effusion for Sodium?
Sodium = Na
Molar mass of Na = 23.0
1/√23.0 = .21

21. Practice Problems
Find the rate of effusion for the following elements:
Lead
Oxygen
Bromine
Copper
Explain the pattern.

22. Effusion Rates
Lighter particles diffuse more rapidly than heavier particles
Different gases at the same temperature should have the same average kinetic energy
Ratio of Diffusion Rates
Ammonia has a molar mass of 17.0 g and HCl has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

23. Practice Problems
Pg 405 #’s 1-3

24. More Practice
Practice
Page 984 #’s 1 and 2

25. Review Go over practice problems pg 405 #’s 1-3
Go over Pg 984 #’s 1 and 2

26. Journal
Why is stepping on a nail more painful than lying on a bed of nails?
Or, why can you ski across deep snow, but not walk across it in boots?

27. Gas Pressure
Pressure
Force per unit area
Because of the small size of an individual gas particle, little pressure is exerted.
When you have many particles, a high pressure can be established.

28. Air Pressure
Because particles in the air move in all directions, pressure is exerted in all directions.
Atmospheric pressure
Pressure is stronger at surface
More particles at surface
More particles above
Stronger pull of gravity

29. Classwork Data Analysis Lab Pg 408 #’s 1-3
Work with shoulder partner!!!!!

30. Understanding Check Read Daltons Law of partial pressure
Review Data Lab Analysis

31. Measuring Air Pressure Evangelista Torricelli
Barometer
Manometer
An instrument used to measure atmospheric Pressure
Higher mercury measurement = higher atmospheric pressure
Determined by:
Gravity exerting a downward force on mercury
Air pressure, exerting a downward force on the dish
Used to measure gas pressure in a closed container

32. Units of Pressure SI unit of pressure = Pascal (Pa)
Many other units used to represent pressure
psi = lbs/in2
mm Hg = millimeter of mercury
torr
bar
Air pressure = atmosphere (atm)
1atm = 760 mm Hg, 760 torr, kPa

33. Daltons Law of Partial Pressure
Total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture
All are independent of one another
Pressure exerted by a single gas = partial pressure
Depends only on moles, size of container, and temperature

34. Calculation Ptotal = P1 + P2 + P3 + …Pn Example:
A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of .97atm, what is the partial pressure of O2, if the partial pressure of CO2 is .7 atm and the partial pressure of N2 is .12?

35. Page 409 #’s 4-7
Page 984 #’s 3 and 4

36. Review Questions
page 410 #’s 8-13

37. Journal
Explain how the addition and removal of energy can cause a phase change.

38. Video: Phases of Matter
Take notes on video
You will be required to write a 1 page summary at the end of the video

39. Phase Changes – Fill in Graph
Requires the addition and removal of energy
How does this lead to a phase change?

40. Melting Heat transfers from higher temp to lower temp
Ice cubes in water
Water is warmer
Energy flows from the water to the ice and melts it.
Absorbing energy allows the hydrogen bonds to break
The amount of energy needed depends on the strength of the forces keeping the particles together.
Melting Point: The temp at which the forces holding its crystal lattice together are broken and it becomes a liquid.
ENERGY ADDED

41. Vaporization What is a vapor?
Vaporization: the process by which a liquid changes to a gas or vapor.
Molecules tend to escape from the surface.
Why?
When vaporization only occurs from the surface its called evaporation.
Figure = vapor pressure
ENERGY ADDED

42. Boiling
Vapor pressure of the liquid equals the external or atmospheric pressure is called the boiling point.
ENERGY ADDED

43. Sublimation Changing directly from a solid phase to a gas phase.
Examples:
Solid Iodine
Solid carbon dioxide (dry ice).
Solid air fresheners
ENERGY ADDED

44. Freezing
As heat is removed from the water, the molecules lose kinetic energy and their velocity decreases.
Less energy, stronger bond
Freezing Point: Temperature at which a liquid is converted to a solid.
ENERGY REMOVED

45. Condensation The process by which a gas or a vapor becomes a liquid.
ENERGY REMOVED

46. Deposition Gas or vapor directly to a solid
In the winter when water vapor forms frost on the window.
Snow
ENERGY REMOVED

47. Heat vs. Temp Change
During a phase change, Temperature does not rise

48. Phase Diagram Two variables control the phase of a substance
Temperature
Pressure
Phase is a description of BOTH temperature and pressure (opposite effects)
Triple point: the temp and pressure at which three phases of a substance can coexist.
All six phases can occur
Critical Point: critical temp and pressure where water cannot exist as a liquid.

49. Other Keys Normal boiling and freezing points are different
CO2 can sublime under normal conditions
Can see in the diagram

50. Practice Problem
You must do the following:
Answer all questions, show me once you receive 100%. This will be a grade.

51. Question of the Day
If all materials at the same temperature have atoms with the same kinetic energy, then why are some solids, some liquids, and others gases?

52. Classwork Watch video on intermolecular forces

53. Intermolecular forces
The attractive forces that hold particles together in ionic, covalent, and metallic bonds are called intramolecular forces.
Do NOT account for all the attraction
Intermolecular forces
Holds together identical particles
Three forces
Dispersion
Dipole-dipole
Hydrogen

54. Dispersion Forces
Weak forces that result from temporary shifts in the density of electrons in electron clouds.
a.k.a London forces
Occurs when molecules come into close contact with one another, and density of electrons change
Increases force as the number of electrons increase

55. Dipole-Dipole Forces
Just like polar molecules where they have a slightly positive charge and a slightly negative charge.
The attraction is dipole-dipole forces
Orient so that oppositely charged particles attract
Permanent = Stronger

56. Hydrogen Bonds A type of Dipole-Dipole attraction
Occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair.
Strongest
Must be bonded to a F, O, or N atom
Electronegativity it high enough
Positive H is attracted to Negative lone pairs

57. Classwork Watch video on intermolecular forces
Work on practice problems
Pg. 435 #’s 75-82
Pg. 434 #’s 51-59

58. Answer the Following Question
NOW, based on what we learned about intermolecular forces, answer the following questions.
How is it possible for water to be a liquid at room temperature while other similar molecules are a gas?

59. Finish – Test Thurs/Fri
Continue Working on Review Practice Problems
Let me know if you have any questions.
Page 430 #’s 27-33
Page 434 #’s 34-41
Pg. 435 #’s 75-82
Pg. 434 #’s 51-59

60. Journal
Explain why substances have different boiling and melting points

61. Questions 1 and 2
The Kinetic Molecular Theory describes the properties of gases in terms of size, motion, and energy. Explain how this theory describes each.
What is density and how do you explain it in terms of particles?

62. Question 3 and 4
Why can air be compressed? How does density relate to change in pressure?
What is the difference between diffusion and effusion? How can more particles move through the air when it is already filled?

63. Question 5 and 6 What is the rate of effusion of oxygen?
Carbon dioxide has a molar mass of 44 g/mol and HCl has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

64. Questions 7 and 8
How do you get a high pressure when representative particles are so small?
Why is pressure stronger at the surface of the Earth?

65. Questions 9 and 10
What are two tools used to measure pressure? What is the SI unit for pressure?
If the total pressure of air is 101.3, then what is the pressure of O2, when the partial pressures of N2 is 78.1, Ar is .97, H2O is 1.28, and CO2 is .05?

66. Questions 11 and 12 How does a substance melt? Boil? Freeze?
What is a vapor? What is the difference between vaporization and evaporation?

67. Questions 13 and 14
What is the difference between sublimation and deposition?
Explain what happens to temperature during a phase change.
Sketch the graph

68. Questions 15 and 16
Where is the triple point located on the graph? What does the triple point indicate?
What is the difference between intermolecular and intramolecular forces?

69. Questions 17 and 18
Explain the differences between dispersion force, dipole-dipole force, and hydrogen bonds.
Which is the weakest force? Strongest force? Why?

70. Journal Look over your answers to the questions from yesterday.
Look in your book and make corrections or add information.

71. Chapter 12 Study Guide Know all parts of the Kinetic Molecular Theory
What is elastic collision?
Understand what causes particles to move faster or slower.
Will smaller or larger molecules flow faster?
How do you know which one is smaller and which one is larger?

72. How does kinetic energy and temperature relate?
Can a gas and a liquid have the same kinetic energy? Explain.
Understand the phase change graph (heat vs. temperature)
Understand the phase change diagram (Temperature vs. Pressure)
Be able to label solid, liquid, and gas.
Find triple point, critical point, melting point, and boiling point.
Explain boiling, freezing, condensation, evaporation, vaporization, sublimation, and deposition.
Relate it to transfer of energy.
How do you calculate partial pressure?

73. Calculate rate of diffusion.
What is normal atmospheric pressure?
How do you calculate average kinetic energy?

74. What is the difference between intermolecular and intramolecular?
Describe dispersion, dipole-dipole, and hydrogen bonding
Which is the strongest?
How do intermolecular forces correlate to the states of matter?
What causes two molecules with similar masses to be in different states of matter when the same temperature is applied?
Calculate the ratio for the rate of diffusion