1. CHAPTER 10 STATES OF MATTER

2. Sections 10.1 – Kinetic Molecular Theory 10.2 – Liquids 10.3 – Solids
10.4 – Changes of State
10.4 – Water

3. 10. 1 Kinetic Molecular Theory
State the kinetic-molecular theory of matter, and describe how it explains certain properties of matter.
List the five assumptions of the kinetic-molecular theory of gases.
Define the terms ideal gas and real gas.
Describe each of the following characteristic properties of gases: expansion, density, fluidity, compressibility, diffusion, and effusion.
Describe the conditions under which a real gas deviates from “ideal” behavior.

4. What is the Kinetic Molecular Theory?
Break it down:
Kinetic: movement
Molecular: particles
Theory: tested ideas
Tested ideas about the
movement of particles!
This theory is used to explain the energy and forces that cause the properties of solids, liquids, and gases.

5. KMT of Gases
Ideal gas: hypothetical gas based on the following five assumptions…
Gases consist of large numbers of tiny particles that are far apart relative to their size.
Most of the volume is empty space
Collisions between gas particles and between particles and container walls are elastic collisions.
elastic collision when there is no net loss of total kinetic energy

6. KMT cont.
Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy, which is energy of motion.
There are no forces of attraction between gas particles.
The temperature of a gas depends on the average kinetic energy of the particles of the gas.
The kinetic energy of any moving object is given by the following equation:

7. Gas Behavior KMT applies only to ideal gasses.
Most gasses behave ideally if pressure is not too high or temperature is not too low.

8. Expansion Gases: do not have a definite shape or a definite volume.
Gas particles move rapidly in all directions (#3) without significant attraction between them (#4).

9. Fluidity
Because the attractive forces between gas particles are insignificant (#4), gas particles glide easily past one another.
Because liquids and gases flow, they are both referred to as fluids.

10. Low Density
The density of a gas at atmospheric pressure is about 1/1000 the density of the same substance in the liquid or solid state.
The reason is that the particles are so much farther apart in the gaseous state (#1).

11. Compressibility
During compression, the gas particles, which are initially very far apart (#1), are crowded closer together.

12. Diffusion
Gases spread out and mix with one another, even without being stirred.
Spontaneous mixing of the particles of two substances caused by their random motion: diffusion.
The random and continuous motion of the gas molecules (#3) carries them throughout the available space.

13. Effusion
Effusion: when gas particles pass through a tiny opening
The rates of effusion of different gases are directly proportional to the velocities of their particles.
Molecules of low mass effuse faster than molecules of high mass.

14. Diffusion vs. Effusion Video

15. A Real Gas
Real gas: does not behave completely according to the assumptions of the kinetic-molecular theory.
Because particles of gases occupy space and exert attractive forces on each other, all real gases deviate to some degree from ideal gas behavior.
The more polar the molecules of a gas are, the more the gas will deviate from ideal gas behavior.
Conditions for a real gas:
high pressures and low temperatures

16. 10.2 Liquids
Describe the motion of particles in liquids and the properties of liquids according to the kinetic- molecular theory.
Discuss the process by which liquids can change into a gas. Define vaporization.
Discuss the process by which liquids can change into a solid. Define freezing.

17. KMT of Liquids Liquids: definite volume and no definite shape
What does this mean about the energy in liquids compared to gases?
The attractive forces between particles in a liquid are more effective than those between particles in a gas.
This is due to intermolecular forces:
dipole-dipole forces
hydrogen bonding
London dispersion forces

18. Fluidity
fluid: a substance that can flow and therefore take the shape of its container.
The particles in a liquid are not bound together in fixed positions. Instead, they move about constantly and slide past each other.

19. Density and Compressibility
At normal atmospheric pressure…
most LIQUIDS are a hundreds times DENSER than in a gaseous state.
LIQUIDS are much LESS COMPRESSIBLE than GASES because liquid particles are more closely packed together.

20. Diffusion
Any liquid gradually diffuses throughout any other liquid in which it can dissolve.
Because the particles are in constant motion.
Diffusion is much slower in liquids than in gases.
Liquids are more tightly packed.
Intermolecular forces slow liquid movement.
As the temperature increases, diffusion occurs more rapidly.

21. Diffusion of Dye through Water

22. Surface Tension
surface tension: a force that pulls adjacent parts of a liquid’s surface together, decreasing surface area to the smallest possible size (cohesion)

23. The higher the intermolecular forces between the particles of a liquid, the higher the surface tension.
The molecules at the surface of the water can form hydrogen bonds with the other water, but not with the molecules in the air above

24. Surface Tension Video

25. Adhesion
Capillary action: attraction of the surface of a liquid to the surface of a solid (adhesion)
This attraction pulls the liquid molecules upward along the surface against the pull of gravity.
This causes the concave liquid surface, called a meniscus, that forms in a test tube or graduated cylinder

26. Capillary Action Video

27. Phase Change to a Gas
Vaporization: process by which a liquid or solid changes to a gas
Evaporation: when particles escape from the surface of a non-boiling liquid and enter the gas state.
Boiling change of a liquid to bubbles of vapor that appear throughout the liquid.
Evaporation occurs because the particles of a liquid have different kinetic energies.

28. Phase Change to a Solid
When a liquid is cooled, the average energy of its particles decreases.
Freezing or solidification: physical change of a liquid to a solid by removal of energy as heat

29. Phase Change Video

30. 10.3 Solids
Describe the motion of particles in solids and the properties of solids according to the kinetic- molecular theory.
Distinguish between the two types of solids.
Describe the different types of crystal symmetry.
Define crystal structure and unit cell.

31. KMT of Solids Solids: have definite shape and definite volume
The particles are tightly packed
All inter-particle attractions exert stronger effects in solids than in the corresponding liquids or gases.
Attractive forces hold the particles in relatively fixed positions.
Solids are more ordered than liquids and are much more ordered than gases.

32. Types of Solids
crystalline solids: the most common type of solid made of crystals
crystal: substance in which the particles are arranged in an orderly, geometric, repeating pattern.
amorphous solid: when the particles are arranged randomly.
Crystalline 
 Amorphous

33. Phase Change of a Crystal
Melting: physical change of a solid to a liquid by the addition of energy as heat.
Which has a higher melting point, ionic or covalent compounds?
The melting point is when…
the kinetic energies of the particles within the solid overcome the attractive forces holding them together.
Sublimation: change from solid directly to a gas

34. Phase Change of Amorphous Solid
Amorphous solids have no definite melting point.
example: glass and plastics
Amorphous solids are sometimes classified as supercooled liquids, which are substances that retain certain liquid properties even at temperatures at which they appear to be solid.
These properties exist because of the arrangement of amorphous solids.

35. Properties of Solids
Substances are usually most dense in the solid state.
Why are solids usually the most dense?
Why is water an exception?
Solids can be considered incompressible.
Solids diffuse millions of times slower than liquids.

36. Crystals
Solids exist as single crystals or as groups of crystals fused together.
The total three-dimensional arrangement of particles of a crystal is called a crystal structure.
The arrangement of particles in the crystal can be represented by a coordinate system called a lattice.
The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice is called a unit cell.

37. Unit Cells

38. Ionic Crystals
Ionic crystals — positive and negative ions arranged in a regular pattern.
Generally, ionic crystals form when Group 1 or Group 2 metals combine with Group 16 or Group 17 nonmetals or nonmetallic polyatomic ions.
These crystals are hard and brittle, have high melting points, and are good insulators.

39. Covalent Network Crystals
Covalent network crystals — each atom is covalently bonded to its nearest neighboring atoms.
The covalent bonding extends throughout a network that includes a very large number of atoms. Represented by a metal and a polyatomic ion.
The network solids are very hard and brittle, have high melting points and are usually nonconductors or semiconductors.

40. Metallic Crystals
Metallic crystals — metal cations surrounded by a sea of delocalized valence electrons.
The electrons come from the metal atoms and belong to the crystal as a whole.
The freedom of these delocalized electrons to move throughout the crystal explains the high electric conductivity of metals.

41. Covalent Molecular Crystals
Covalent molecular crystals — covalently bonded molecules held together by intermolecular forces.
Nonpolar molecules: weak London dispersion forces between molecules
Polar molecules: dispersion forces, dipole-dipole forces, and sometimes hydrogen bonding
Covalent molecular crystals have low melting points, are easily vaporized, are relatively soft, and are good insulators.

42. Phase Change with Heating Curve

43. 10.4 Change of State
Explain the relationship between equilibrium and changes of state.
Interpret phase diagrams.
Explain what is meant by equilibrium vapor pressure.
Describe the processes of boiling, freezing, melting, and sublimation.

44. Phase Changes

45. Key Words
Phase: part of a system that has uniform composition and properties.
Condensation: gas changes to a liquid.
Vapor: gas in contact with its liquid or solid phase

46. Equilibrium
Equilibrium: condition in which two opposing changes occur at equal rates in a closed system.
For example:
In a closed system, the rate of condensation equals the rate of evaporation, and a state of equilibrium is established.

47. Equilibrium of Liquid and Vapor

48. Vapor Pressure
equilibrium vapor pressure: pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature
The equilibrium vapor pressure increases with increasing temperature.
Increasing the temperature of a liquid increases the average kinetic energy of the liquid’s molecules.
Every liquid has a specific equilibrium vapor pressure at a given temperature.

49. Equilibrium Vapor Pressure

50. Equilibrium Vapor Pressure Video

51. Volatile Liquids Volatile liquids: liquids that evaporate readily.
They have relatively weak forces of attraction between their particles.
example: ether
Nonvolatile liquids: do not evaporate readily.
They have relatively strong attractive forces between their particles.
example: molten ionic compounds

52. Which substance is most volatile?

53. Volatile Liquids Video

54. Boiling
Boiling Point: temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure.
The lower the atm pressure is, the lower the boiling point is.
At the boiling point, all of the energy absorbed is used to evaporate the liquid, and the temperature remains constant as long as the pressure does not change.
If the pressure is increased, the temperature of the liquid will rise until the vapor pressure equals the new pressure and the liquid boils once again.

55. The normal boiling point of a liquid is the boiling point at normal atmospheric pressure (1 atm, 760 torr, or kPa).
The normal boiling point of water is exactly 100°C.

56. Molar Enthalpy of Vaporization (L/G)
Energy must be added continuously in order to keep a liquid boiling.
molar enthalpy of vaporization ∆Hv
amount of energy as heat that is needed to vaporize one mole of liquid at the liquid’s boiling point at constant pressure
The stronger this attraction is between liquid particles, the higher molar enthalpy of vaporization.

57. ∆Hv Calculation

58. Each liquid has a characteristic molar enthalpy of vaporization.
Why does water have an unusually high molar enthalpy of vaporization ??

59. Freezing liquid solid + energy Freezing: liquid to a solid
loss of energy in the form of heat by the liquid.
In the case of a pure crystalline substance, this change occurs at constant temperature.
liquid solid + energy

60. Freezing cont. solid + energy liquid
normal freezing point: temperature at which the solid and liquid are in equilibrium at 1 atm (760 torr, or kPa) pressure.
Melting occurs at constant temperature, both the solid and liquid particles have the same average kinetic energy.
solid + energy liquid

61. Phase Change Equilibrium
At normal atmospheric pressure, the temperature of a system containing ice and liquid water will remain at 0.°C as long as both ice and water are present.
After all the ice has melted will the addition of energy increase the temperature of the system!

62. Molar Enthalpy of Fusion (S/L)
molar enthalpy of fusion ∆Hf
amount of energy as heat required to melt one mole of solid at the solid’s melting point is the solid’s
The magnitude of the molar enthalpy of fusion depends on the attraction between the solid particles.

63. ∆Hf Calculation

64. Phase Change Between Solid and Gas
At low pressures liquids can not exist
Sublimation: solid to a gas
Deposition: gas to a solid

65. Phase Diagrams
phase diagram: graph of pressure vs temperature that shows the conditions under which the phases of a substance exist.
triple point: temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium.

66. Critical Point
critical point of a substance indicates the critical temperature and critical pressure.
critical temperature (Tc) temperature above which the substance cannot exist in the liquid state.
Above this temperature, a substance becomes a supercritical fluid
critical pressure (Pc ) lowest pressure at which the substance can exist as a liquid at the critical temperature.

67. Phase Diagram of Water Where is the critical point?
Where is the triple point?
At what pressure is the normal freezing and boiling point?
Where does a substance become a supercritical fluid?
What is the density relation between solid and liquid?

68. Phase Diagram of Carbon Dioxide
Where is the critical point?
Where is the triple point?
At what pressure is the normal freezing and boiling point?
Where does a substance become a supercritical fluid?
What is the density relation between solid and liquid?

69. Phase Changes Review

70. Heating Curve of Water

71. What would a Cooling Curve look like?
Energy released

72. 10.5 Water Describe the structure of a water molecule.
Discuss the physical properties of water. Explain how they are determined by the structure of water.
Calculate the amount of energy absorbed or released when a quantity of water changes state.

73. Structure of Water What is the formula for water?
What type of bond connects the hydrogen and oxygen?
What type of intermolecular force occurs between water molecules?
The number of linked molecules decreases with increasing temperature.
Ice consists of water molecules in a crystal structure.

74. Molecules of Water

75. Which is more dense?

76. The hydrogen bonds between molecules of liquid water at 0°C are fewer and more disordered than those between molecules of ice at the same temperature.
Liquid water is denser than ice.
As the temperature approaches the boiling point, groups of liquid water molecules absorb enough energy to break up into separate molecules.

77. Remember the Phase Diagram

78. Physical Properties
At room temperature, pure liquid water is transparent, odorless, tasteless, and almost colorless.
The molar enthalpy of fusion of ice is relatively large compared with the molar enthalpy of fusion of other solids. Why?
Water expands in volume as it freezes, because its molecules form an open rigid structure.
This lower density explains why ice floats in liquid water.

79. Steam (vaporized water) stores a great deal of energy as heat.
Both the boiling point and the molar enthalpy of vaporization of water are high compared with those of nonpolar substances of comparable molecular mass.
The values are high because of the strong hydrogen bonding that must be overcome for boiling to occur.
Steam (vaporized water) stores a great deal of energy as heat.