1. Kinetic Molecular Theory

2. Kinetic theory Describes how particles move as an ideal gas
Can be adjusted to describe the motion of particles in all states of matter
Points of the Kinetic Theory
Gases consist of a large number of tiny particles that are far apart relative to their size
Collisions between particles and between particles and the container wall are elastic
Gas particles are in continuous, rapid, random motion
There are no forces of attraction between gas particles
The temperature of a gas depends on the average kinetic energy of the particles in the gas
KE = 1/2mv2

3. Behavior of Gases
The behavior of a gas is directly related to the five points of the theory
Gases Expand (expansion)
This is a result of no attractive forces and a constant, random motion
Gases are fluids (fluidity)
Particles can freely move past one another
No attractive forces, tiny particles that are far apart, and constant motion
Gases have a low density
A lot of empty space with a relatively low mass
Particles are far apart relative to their size

4. Behavior Continued Gases can be compressed
Gases are far apart with empty space therefore they can be forced closer
Gases have the ability to diffuse and effuse
Due to the constant, random motion of gases particles, they will fill the volume of their container and no attractive forces they will mix together
This allows them to mix evenly (diffusion) and escape through tiny openings in their containers (effusion)

5. Real Gases
Real gases tend to deviate from ideal gases in that we do see some slightly attractive forces between particles
Gases such as the noble gases will exhibit a more ideal nature than other gases due to their unreactive nature
The temperature and pressure can also affect the ideal nature
High temperatures and low pressures will typically exhibit an ideal nature

6. Liquids
As temperature and pressure begin to change, the principles of the kinetic theory begin to breakdown
Principle #4: There are no attractive forces between particles
As kinetic energy decreases, attraction between particles will increase
Particles in the liquid state lose kinetic energy and become attracted to each other which forces the particles to stay close together (fixed volume)

7. Behavior of Liquids Liquids are fluids Liquids have a high density
Even though they lose kinetic energy, particles still have enough kinetic energy to rotate freely around each other
Liquids have a high density
With stronger attractive forces, particles are more tightly packet occupying a smaller volume
Less empty space between particles
Liquids are relatively incompressible
Eliminate empty space
Liquids can diffuse
Particles can still move freely around one another so they can freely mix with other liquids

8. Behaviors Continued Liquids have surface tension
Due to the attractive forces, liquid particles want to stay together so there is a force that is required to break them apart
There is also an attraction between the liquid particles and solid particles known as capillary action
Liquids have the ability to evaporate and boil
Both boiling and evaporation are part of a process known as vaporization (liquid becoming a gas)
Evaporation occurs only at the surface of the liquid
Boiling can occur throughout the liquid
Liquids can become solids by freezing

9. Solids
With a continued loss of kinetic energy, the attractive forces between particles will continue to grow stronger pulling the particles closer together
Eventually, the particles will be pulled into a fixed position and exhibit a repeating pattern
Particles will just be vibrating in place

10. Behavior of Solids Solids have two types: crystalline and amorphous
Crystalline solids consist of crystals which have an orderly, geometric, repeating pattern
Amorphous solids have a random arrangement of particles
Solids have a definite shape and volume
Solids have a definite melting point
Defined temperature required to break apart the attractive forces and allow the particles to move freely
Solids have a high density and are incompressible
Solids have a very low rate of diffusion
The rate is millions of times slower in solids than in liquids

11. Crystalline Solids Crystal structure
Three dimensional arrangement of particles
This arrangement can be represented by a coordinate system called a lattice
A unit cell is the smallest representation of this repeating pattern

12. Forces of attraction Crystals can be classified into four categories
Ionic Crystals: repeating pattern of oppositely charged ions
Covalent Network Crystals: each atom is bonded to the nearest atom by sharing electrons
Metallic Crystals: Repeating pattern of metal cations that are surrounded by a sea of electrons
Covalent Molecular crystals: Molecules are formed by covalently bonded atoms and then the molecules are held together by intermolecular attraction

13. Amorphous Solids Comes from the Greek word meaning “without shape”
Glasses and Plastics (polymers) are good examples
These solids are formed in a manner that does not allow them to crystallize whether it be by a cooling process (glasses) or molded under high temperature and pressure (plastics)

14. Changes of State Possible changes in state Solid to liquid (melting)
Solid to gas (sublimation)
Liquid to solid (freezing)
Liquid to gas (vaporization)
Gas to liquid (condensation)
Gas to solid (deposition)

15. Equilibrium
This is a dynamic condition in which two opposing changes occur at equal rates in a closed system
Equilibrium can be establish between any two states that are changing at the temperature where that change occurs
For example, a sealed liquid will continually build up vapor pressure at the surface of the liquid as it evaporates
This process will continue until equilibrium vapor pressure was achieved
Keeping in mind that equilibrium is a dynamic condition
Temperature and pressure can change the equilibrium points

16. Boiling
Boiling occurs when the partial pressure of the liquid is equal to the atmospheric pressure
Boiling point is the temperature where boiling will occur
Normal boiling point is the temperature at 1 atm of pressure (or 760 torr or kPa)
Once a substance has started to boil, all the energy that is added goes directly to converting the state
The amount of energy needed to vaporize one mole of liquid at the boiling point is known as the molar enthalpy (heat) of vaporization
Expressed as kJ/mol (kilojoules per mole)

17. Freezing and Melting
Freezing point is the temperature at which a solid and liquid are in equilibrium
Normal freezing point is also at 1atm
Melting point is the same temperature just depends on whether energy is added or removed from our system
Just like with boiling, energy will be added (melting) or removed (freezing) until the substance is entirely converted but the temperature will not change
Water at 0ºC will remain at 0ºC until all the ice is melted or the water is frozen
The amount of energy required to melt one mole of a solid at the melting point is the molar enthalpy (heat) of fusion
Expressed as kJ/mol

18. Phase Diagram

19. Phase Diagram Triple Point Critical Point Critical temperature
Temperature and pressure where all three state exist in equilibrium
Critical Point
Indicates where the critical temperature and critical pressure meet
Critical temperature
Temperature in which you can no longer have a substance in the liquid state
Critical pressure
Lowest pressure where a substance can still be a liquid at the critical temperature