1. Chapter 13 States of Matter

2. FOOD FOR THOUGHT…
What is the relationship between solids, liquids, and gases? How are they the same? How are they different?
Make a table or list answering the above question. You will need your answers tomorrow.

3. Gases
Flemish physician Jan Baptista Van Helmont used the Greek word chaos, meaning without order, to describe the products or reactions that had no fixed shape or volume.
From chaos came the term gas.

4. Gases
Around 1860, Ludwig Boltzmann and James Maxwell each proposed a model to explain the properties of gases. The model is known as the kinetic-molecular theory.
Kinetic-molecular theory describes the behavior of gases in terms of particles in motion. 4

5. 13.1 - Gases Kinetic – molecular theory: (pages 386, 419-420)
Gas particles do not attract or repel each other.
Why?
Gas particles are much smaller than the distance between them. The theory assumes the particles have virtually no volume.
What process can be applied to gases because of the very small particle size?
Gas particles are in constant, random motion. Because of this, gas particles spread out and mix.
How do gas particles move? What may cause their motion to change? 5

6. 13.1 - Gases Kinetic – molecular theory: (pages 386, 419-420)
No kinetic energy is lost when gas particles collide with each other or with the walls of their container.
Why?
How can you describe elastic collisions?
What can be said about the total kinetic energy of a container of gas?
All gases have the same average kinetic energy at a given temperature.
What happens to total energy of a gas as temperature increases?
What happens to total energy of a gas as temperature decreases?
What two factors determine kinetic energy?
What is temperature? What do all gases have in common? 6

7. Food for thought…
Will the kinetic-molecular theory work for liquids and solids?
Explain your reasoning.

8. Gases
How does kinetic-molecular theory explain the behavior of gases?
Density
The idea that there is a lot space between particles explains the low density (mass / volume) that gases have. There are fewer molecules of gas in a given amount of volume.
Compression and Expansion
Gas molecules can be compressed (pressed into a smaller space using a piston).
Gas molecules can expand as pressure from the piston is released.
What happens to the density as a gas is compressed?
What happens to the density as a gas expands? 8

9. Gases
How does kinetic-molecular theory explain the behavior of gases?
Diffusion
The idea that there are no significant forces of attraction between gas particles supports the fact that gases can flow easily past each other.
When gases mix, what will eventually happen to their concentration?
What is diffusion?
In which direction does diffusion occur?
What does the rate of diffusion depend upon?
Effusion
What is effusion?
How can effusion be compared to diffusion? 9

10. Gases
How does kinetic-molecular theory explain the behavior of gases?
Effusion
Graham’s law of effusion – the rate of effusion for a gas is inversely proportional to the square root of its molar mass
Graham’s law applies to diffusion rates also.
Rate of effusion
SQRT(molar mass)
Rate A = SQRT(molar mass B / molar mass A)
Rate B 10

11. Food for thought… What state of matter is He?
What is happening to the He atoms inside a balloon to keep the balloon inflated? Use the kinetic-molecular theory to support your answers.

12. Gases
TIME TO WORK!
Effusion / diffusion problems (1-2), page 388
Dalton’s Law problems (4-6), page 392
Problem-solving lab, page 390
Calculate the ratio of effusion rates for pairs of the noble gases. 12

13. 13.2 – Forces of Attraction Intramolecular Forces
The forces that hold particles together in ionic, covalent, and metallic bonds are called intramolecular forces. These are forces occurring within the chemical compound.
Intermolecular Forces
Intermolecular forces happen between or among like molecules of a substance.
There are 3 intermolecular forces we will discuss: dispersion forces, dipole-dipole forces, and hydrogen bonds. 13

14. 13.2 – Forces of Attraction Intramolecular Forces
All intermolecular forces are weaker than intramolecular bonding forces.
Relative Strength of Molecular Forces
Covalent network > ionic bonds > metallic bonds >
Hydrogen bonds > dipole-dipole forces > London dispersion
forces 14

15. Food for thought…
Many words can be understood by looking at their parts.
Match the following word parts to their meanings.
-ion therm-
-ic ize
endo-
exo-
Heat
The result of an action
Related to
To become
Inside
Outside

16. 13.2 – Forces of Attraction London dispersion forces –
are weak forces that result from temporary shifts in the density of electrons in electron clouds (draw sketch page 394)
are named after the German-American physicist who first described them, Fritz London
are weak forces because they are based on temporary dipoles
are the dominant force of attraction between identical nonpolar molecules
dispersion force strength: I > Br > Cl > F
explains why F and Cl are gases, Br is liquid, and I is solid at room temperature 16

17. 13.2 – Forces of Attraction Dipole-dipole forces –
are forces of attraction between oppositely charged regions of polar molecules.
are weak forces of attraction that result from permanent dipoles within polar molecules
Some regions of a polar molecule are always partially negative and other regions are always partially positive.
Neighboring polar molecules orient themselves so that oppositely charged regions line up. (sketch diagram page 394)
The degree of polarity in a molecule depends on the relative electronegativity values of the elements in the molecule.
Dipole forces are stronger than dispersion forces as long as the molecules being compared have about the same mass. 17

18. Bell ringer…
Which are stronger, intermolecular or intramolecular forces?
Which specific force is strongest?
Which is weakest?

19. 13.2 – Forces of Attraction Hydrogen bonds –
are a special type of dipole attractive force between highly polar molecules.
are dipole-dipole attractions that occur between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair. (sketch diagram page 395)
Hydrogen must be bonded to either a F, O, or N atom because these atoms are electronegative enough to cause a large partial positive charge on the hydrogen, yet are small enough that their lone pairs of electrons can come close to hydrogen atoms.
In hydrogen bonds, hydrogen atoms on one molecule of a substance are attracted to the negative end of another molecule of the substance.
Hydrogen bonds explain why water is liquid at room temp. 19

20. 13.3 – Solids and Liquids
Although kinetic-molecular theory was developed to explain the behavior of gases, the model can be applied to liquids and solids. However, you must consider the forces of attraction between particles as well as their energy of motion.
Why consider the force of attraction between particles of liquids and solids?
In liquids and solids, particles are closer together because of stronger forces of attraction between them. In liquids, the forces of attraction limit their range of motion so that particles are closely packed in a fixed volume.
In solids, the forces are so strong that motion is limited to vibrations around a fixed location. Therefore, solids have more order in their particles. 20

21. Food for thought…
Put the following attractive forces in the proper category: intermolecular or intramolecular
dispersion forces dipole-dipole covalent bonds
metallic bonds hydrogen bonds ionic bonds
Put the attractive forces in order from weakest to strongest.

22. 13.3 – Solids and Liquids
Liquids - define the following terms related to properties of liquids
Density
Compression
Fluidity
Viscosity
Temperature (relationship to viscosity)
Surface tension (define and give an example)
Surfactants
Capillary action (define and give an example) 22

23. 13.3 – Solids and Liquids
Solids - define the following terms related to properties of solids
Density
Compressibility
Fluidity
Crystalline solid
Unit cell
Molecular solids
Covalent network solids
Ionic solids
Metallic solids
Amorphous solids 23

24. 13.3 – Solids and Liquids
Sketch and label 3 basic solid crystal lattice structures (page 400).

25. 13.4 – Phase Changes 6 possible transitions between phases GAS
SUBLIMATION
CONDENSATION
DEPOSITION
VAPORIZATION
LIQUID
SOLID
MELTING
FREEZING 25

26. 13.4 – ENDOTHERMIC CHANGES Phase changes requiring energy input
Melting
What is melting?
What is heat?
Describe what happens when ice melts.
Why is the energy required to melt salt much greater than the energy to melt ice?
What is melting point? 26

27. Endothermic Phase Changes
Phase changes requiring energy input
Melting
Melting is when energy added to a solid is great enough to
break the forces holding the molecules together.
The solid phase changes to a liquid.
Melting point of a solid is the temperature at which forces holding its structure together are broken and it becomes a liquid.
Heat is the transfer of energy from a higher temperature to a
lower temperature.
When ice melts, molecules on the surface absorb enough
Energy to break the hydrogen bonds. The molecules move
apart and become a liquid.
The ionic bonds in salt (sodium chloride) are much stronger
than the hydrogen bonds in ice. 27

28. 13.4 – ENDOTHERMIC CHANGES Phase changes requiring energy input
Vaporization
When can the temperature of a melting substance start to rise?
How much energy does a substance need to vaporize?
What is a vapor?
What is vaporization?
What is evaporation?
How long does it take for evaporation to occur? 28

29. Endothermic Phase Changes
Phase changes requiring energy input
Vaporization
The temperature of a melting substance will start to rise after
all the solid substance has melted.
Evaporation is vaporization that occurs only at the surface of a liquid.
To vaporize, a substance needs enough energy to overcome
the forces of attraction holding the molecules together in the
liquid.
The time for evaporation depends upon the amount of liquid and the amount of energy available.
A vapor is a substance that is ordinarily a liquid at room
temperature.
Vaporization is the process by which a liquid changes to a gas
or vapor. 29

30. 13.4 – ENDOTHERMIC CHANGES Phase changes requiring energy input
Vaporization
What is vapor pressure?
How does a rise in temperature affect vapor pressure?
What is the boiling point?
As temperature increases, what happens to kinetic energy of molecules? 30

31. Endothermic Phase Changes
Phase changes requiring energy input
Vaporization
Vapor pressure is the pressure exerted by a vapor over a
liquid.
Vapor pressure increases with increasing temperature.
The boiling point is the temperature at which the vapor
pressure of a liquid equals the outside or atmospheric
pressure.
Kinetic energy of molecules increases with temperature. 31

32. Endothermic Phase Changes
Phase changes requiring energy input
Sublimation
What is sublimation?
Name some substances that sublime.
Sublimation is the process by which a solid changes directly
to a gas without becoming a liquid.
Solid iodine, solid carbon dioxide (dry ice), moth balls,
air fresheners, ice cubes left in freezer. 32

33. 13.4 – Exothermic Changes Phase changes that release energy
Condensation
Name some examples of condensation.
What is condensation?
Describe what happens to molecules when they condense.
What happens when hydrogen bonds form in liquid water?
Name some circumstances that will cause condensation. 33

34. Exothermic Phase Changes
Phase changes that release energy
Condensation
Examples: dew, frost, water on outside of a glass, water on
a window pane, clouds, fog, air conditioning condensate
Condensation can be caused by vapor molecules contacting a cold surface (drink glass, dew) or cold air (fog, clouds).
Condensation is the process by which a gas or a vapor
becomes a liquid. It is the reverse of vaporization.
When vapor molecules condense, they lose energy, slow down, and form bonds with each other when they collide. The bonded molecules are more dense and become a liquid.
When hydrogen bonds are formed in water, energy is released. 34

35. Exothermic Phase Changes
Phase changes that release energy
Deposition
What is deposition?
Give some examples of deposition.
Deposition is the process by which a substances changes
from a gas or vapor to a solid without first becoming a liquid.
It is the reverse of sublimation.
Example: snow 35

36. 13.4 – Exothermic Changes Phase changes that release energy Freezing
What is freezing?
Describe what happens when something freezes.
What is freezing point? 36

37. Exothermic Phase Changes
Phase changes that release energy
Freezing
Freezing is the phase change of a liquid to a solid.
During freezing, heat is removed from a substance, the
molecules lose kinetic energy and slow down. When enough
energy has been removed, the molecules become fixed in a
set position.
The freezing point is the temperature at which a liquid is
converted into a crystalline solid. The freezing point and
melting point temperatures are the same for a substance. 37

38. 15. SUMMARIZING ________thermic _______  Solid Gas  ______
Liquid  ______
Order of the molecules is
__________________.
________thermic
_______  Solid
Gas  ______
______  Solid
Order of the molecules is
__________________.

39. SUMMARIZING gaining ENDOTHERMIC
Molecules are_________ energy during an endothermic phase change.
What kind of energy?
Describe the change in motion of molecules during melting, vaporization, and sublimation.
gaining
Kinetic energy – the energy of motion
The molecules gain energy and move faster and further apart.

40. SUMMARIZING losing EXOTHERMIC
Molecules are__________ kinetic energy during an exothermic phase change.
Describe the change in motion of molecules during freezing, condensation, and deposition.
losing
The molecules lose energy and slow down getting closer together.

41. 13.4 – Phase Changes Phase diagrams
Temperature and pressure are the two variables that combine to control the phase of a substance.
A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure.
What is the triple point?
All 6 phase changes can occur at the triple point. 41

42. 13.4 – Phase Changes
Phase diagrams
What is the critical point? 42

43. Problem solving! Add to your tools!
Read the problem statement thoroughly.
Identify the “givens” and write them down.
Identify the “unknowns” and write them down.
Determine the scientific principles you will use.
Devise a strategy for attacking the problem.

44. Scientific Principles
Name principles you might use to solve various types of physical science problems.
Law of Conservation of Mass
Rules for Balancing Formulas and Equations
Periodic Table
Solving Mathematical Equations, like F = ma, velocity, acceleration, momentum
Rules for Balanced and Unbalanced Forces
Gravity – free fall, weight, acceleration due to gravity
Newton’s Laws

45. Strategies you have used to solve problems…
Using graphs to show data.
Drawing diagrams to show atomic structure.
Using the Periodic Table to determine trends in atomic radius and ionization energy.
Balancing chemical equations.
Using equations to calculate an unknown value.
Drawing a force diagram and showing the magnitude and direction of the forces acting on an object.