1. Chapter 10 States of Matter
10.1 The Kinetic-Molecular Theory of Matter

2. Kinetic-Molecular Theory of Gases
Particles of matter are ALWAYS in motion; constant, rapid motion. (kinetic energy!)
Particles are very small & relatively far apart.
Collisions of particles with container walls cause pressure exerted by gas.
Volume of individual particles is  zero.
Particles exert no attractive or repulsive forces on each other.

3. Kinetic-Molecular Theory of Gases
Gas particles undergo elastic collisions: Collisions in which no energy is lost
Average kinetic energy is directly proportional to Kelvin temperature of a gas.
Air Hockey Table

4. Ideal Gas
An imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory
A gas with its particles in constant random motion without attraction for each other is called an Ideal Gas. These particles undergo elastic collisions.
Nearly all real gases behave as ideal gases EXCEPT at very low temperatures or high pressures.

5. Likely to behave nearly ideally: Likely not to behave ideally:
Real Gases
A gas that does not behave completely according to the assumptions of the kinetic-molecular theory.
Real gases occupy space and exert attractive forces on one another
Likely to behave nearly ideally:
Likely not to behave ideally:
high temp. & low pressure
low temp. & high pressure
Small non-polar gas molecules
Large polar gas molecules

6. Kinetic-Molecular Theory of the Nature of Gases
Expansion
Gases do not have a definite shape or volume
Gases take the shape of their containers
Gases evenly distribute themselves within a container
Fluidity
Gas particles easily flow past one another
Low Density
A substance in the gaseous state has 1/1000 the density of the same substance in the liquid or solid state
Compressibility
Gases can be compressed, decreasing the distance between particles, and decreasing the volume occupied by the gas

7. Kinetic-Molecular Theory of the Nature of Gases
Diffusion
Spontaneous mixing of particles of two substances caused by their random motion
Rate of diffusion is dependent upon:
speed of particles
diameter of particles
attractive forces between the particles

8. Kinetic-Molecular Theory of the Nature of Gases
Effusion
Process by which particles under pressure pass through a tiny opening
Rate of effusion is dependent upon:
speed of particles (small molecules have greater speed than large molecules at the same temperature, so the effuse more rapidly)

9. Chapter 10 States of Matter
10.2 Liquids

10. Some Properties of a Liquid
Surface Tension:
The resistance to an increase in its surface area (polar molecules, liquid metals).
A force that tends to pull adjacent parts of a liquid's surface together, thereby decreasing surface area to the smallest possible size.

11. Some Properties of a Liquid
Capillary Action: Spontaneous rising of a liquid in a narrow tube.

12. Some Properties of a Liquid
Viscosity: Resistance to flow (molecules with large intermolecular forces).

13. Some Properties of Liquids
Volatility
Liquids that have weak forces of attraction and evaporate easily
Nonvolatile Liquids
Liquids that have strong forces of attraction and do not evaporate easily

14. Properties of Fluids Relative High Density
10% less dense than solids (average)
Water is an exception
1000x more dense than gases
Relative Incompressibility
The volume of liquids doesn't change appreciably when pressure is applied
Ability to Diffuse
Liquids diffuse and mix with other liquids
Rate of diffusion increases with temperature

15. Chapter 10 States of Matter
10.3 Solids

16. Types of Solids
Crystalline Solids: highly regular arrangement of their components
[table salt (NaCl), pyrite (FeS2)].

17. Types of Solids
Amorphous solids aka supercooled liquids: considerable disorder in their structures (glass).
Greek for "without shape"
Formation of amorphous solids:
Rapid cooling of molten materials can prevent the formation of crystals
* They do not have definite melting points

18. Representation of Components in a Crystalline Solid
Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

19. Types of Crystalline Solids
Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

20. The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice
Unit Cell

21. Types of Crystalline Solids
Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).

22. Packing in Metals
Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

23. Closest Packing Holes

24. Metal Alloys
Substitutional Alloy: some metal atoms replaced by others of similar size.
brass = Cu/Zn

25. Metal Alloys (continued)
Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.
steel = iron + carbon

26. graphite, diamond, ceramics, glass
Network Solids
Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.
brittle (non-flexible)
do not conduct heat or electricity
carbon, silicon-based
graphite, diamond, ceramics, glass

27. Sulfur – S8

28. Phosphorus – P4

29. Diamond

30. Graphite

31. Zirconia

32. Chapter 10 States of Matter
10.4 Changes of State

33. Equilibrium
Dynamic condition in which two opposing changes occur at equal rates in a closed system
A closed system at constant temperature will reach an equilibrium position at which the rates of evaporation and condensation will be the same

34. Equilibrium Vapor Pressure
The pressure of the vapor present at equilibrium.
Determined principally by the size of the intermolecular forces in the liquid.
Increases significantly with temperature.
Volatile liquids have high vapor pressures.
Increasing the temperature will move more particles into the vapor phase to compensate for the new energy

35. Boiling
The conversion of a liquid to a vapor within the liquid as well as at its surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure
Boiling Point
The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure
Water boils at 100 °C at 1 atm pressure
Water boils above 100 °C at higher pressures
Water boils below 100 °C at lower pressures

36. LeChatelier’s Principle
When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.

37. Translation:
When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away.
When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

38. LeChatelier’s Example #1
A closed container of ice and water at equilibrium. The temperature is raised.
Ice + Energy  Water
The equilibrium of the system shifts to the _______ to use up the added energy.
right

39. LeChatelier’s Example #2
A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container.
N2O4 + Energy  2 NO2
The equilibrium of the system shifts to the _______ to use up the added NO2.
left

40. LeChatelier’s Example #3
A closed container of water and its vapor at equilibrium. Vapor is removed from the system.
water + Energy  vapor
The equilibrium of the system shifts to the _______ to produce more vapor.
right

41. Water phase changes constant Temperature remains __________
during a phase change.
constant

42. Phase Diagram
Represents phases as a function of temperature and pressure.
Critical temperature: temperature above which the vapor can not be liquefied.
Critical pressure: pressure required to liquefy AT the critical temperature.
Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

43. Water
Water

44. Phase changes by Name

45. Carbon dioxide
Carbon
dioxide

46. Carbon
Carbon

47. Sulfur

48. Chapter 10 States of Matter
10.5 Water

49. Water’s Properties

50. Sea Ice
Ice forms on top of the ocean in a thin layer & acts to insulate the warmer waters below from the colder air temperatures.
This occurs in the polar regions, the Artic & Antarctic
Since ice is less dense than liquid water, it will float on top, instead of sinking which would kill all life below the surface.
Sea ice is not the same as an iceberg. Icebergs are pieces of glaciers which are formed by snowfall on land.
Sea ice is not salty, as the hydrogen bonds that hold ice together will not form properly if salt remains in the structure.

51. Temperature & Salinity
Sea Water Density
There are 2 main factors that affect the density of sea water:
Temperature & Salinity
As the temperature decreases, density increases.
As salinity increases, density increases.
Since the least dense layer of a liquid will float above a more dense layer, the warmest & lowest salinity layer will be on top.
However, temperature has a greater effect on density than salinity.
That means that a higher salinity layer can float on top of a lower salinity layer if it is considerably warmer in temperature.

52. Molar Heat of Fusion
The amount of heat energy required to melt one mole of solid at its melting point.
6.009 kJ per 1 mole
Practice!
How much energy is absorbed when 16.3 g of ice melts?
16.3 g x 1 mole x kJ = 5.44 kJ
18 g mol

53. Molar Heat of Vaporization
The amount of heat energy required to vaporize one mole of a liquid at its boiling point
Strong attractive forces between particles result in high molar heat of vaporization
40.79 kJ per 1 mole
Practice!
Find the mass of liquid water required to absorb 5.23 x 104 kJ of energy upon boiling.
5.23 x 104 kJ x 1 mole x 18 g = 2.31 x 104 g kJ 1 mol

54. Specific Heat
The amount of energy required to change the temperature of 1 gram of a substance by 1 degree Celsius is known as specific heat capacity. (Ch.16 p.533)
Q = s * m * ΔT
Where Q = energy (heat) required
s = specific heat capacity
m = mass of sample, grams
ΔT = change in temperature, Celsius

55. Specific Heat Capacities of Some Common Substances
Specific Heat Capacity (J/g°C)
Water, Liquid
4.184
Water, Ice
2.03
Water, Steam
2.0
Aluminum, s
0.89
Iron, s
0.45
Mercury, l
0.14
Carbon, s
0.71
Silver, s
0.24
Gold, s
0.13

56. Practice!
A piece of copper alloy with a mass of 85.0 g is heated from 30.°C to 45°C. In the process, it absorbs 523 J of energy as heat.
What is the specific heat for this copper alloy?
How much energy will the same sample lose if it is cooled from 45°C to 25°C?
Answers…
0.41 J/g∙K
7.0 x 102 J

57. Specific Gravity A function of density.
The ratio of density of a material to the density of water at a specified temperature.
The density of water is usually 1 g/cm3
When dividing density by density, the units cancel and therefore, specific gravity has no units!
If the density of the material is in the units g/cm3 (or g/mL) when dividing by the density of water, the numeric value of the material’s density is the same, but the units cancel out.

58. Intermolecular Forces (Review from Chapter 6)
Forces of attraction between different molecules rather than bonding forces within the same molecule.
Dipole-dipole attraction
Hydrogen bonds
Dispersion forces

59. Dipole-Dipole Attraction
dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “-” ends of the dipoles are close to each other.

60. Dipole Forces

61. Hydrogen Bonding
hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

62. Hydrogen Bonding in Water

63. Hydrogen Bonding in DNA

64. London Dispersion Forces
Dispersion forces: relatively weak forces caused by instantaneous dipole, in which electron distribution becomes asymmetrical.
The ONLY forces of attraction that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18)

65. London Dispersion Forces

66. Boiling point as a measure of intermolecular attractive forces