1. States of Matter There are three main states of matter
Kinetic molecular Theory (KMT) helps us explain the states of matter.

2. Kinetic Molecular Theory (KMT)
KMT states that particles of matter are in constant motion.
Molecular motion is defined by:
Kinetic Energy = ½ mv2
Kinetic energy is a reflection of the substances temperature.
(increase temp = increase kinetic energy)
Some assumptions about KMT as they apply to ideal gases:
Gas particles are spaced far apart
Gas particles collide with “elastic” collisions
Particles are in constant, rapid and random motion
Gases are not attracted to each other.
The temperature of a gas reflects the kinetic energy of the gas particles.

3. States of matter Properties of Gases:
Expansion- gases have no definite shape or volume.
Fluidity- gas particles will easily glide past one another
Low density- gases have 1/1000 the density of other substances
Compressible- gas particles can be pushed closer together
Diffusion- gases can easily disperse and mix in space.
Effusion- gases can pass through a tiny opening.

4. Polar, strong attractions
Deviation from ideal
Comparison of gases:
Gas
Deviation?
Why? N2 No
Non-polar diatomic H2 Ne
Noble gas He
NH3
Yes
Polar, strong attractions
H2O(g)
yes

5. Effusion and Diffusion:
Effusion- gas particles escape through a small opening.
example of effusion
Diffusion-ability of molecule to mix.

6. States of Matter: Properties of Liquids:
Liquids are more dense than gases due to attraction of particles.
Liquids are much less compressible than gases, less empty space.
Liquids have fluidity- able to glide past each other and diffuse.
Liquids can boil and evaporate.
Liquids can be formed into solids.

7. States of Matter Properties of Solids:
Definite shape and volume-organized arrangement of particles
Definite melting point- energy requirement to overcome forces holding the particles together
Greatest density/ incompressible- little/no empty space
Low rate of diffusion- fixed position of particles will not allow them to glide and mix.
Organized/crystalline arrangement
Simulation and states of matter.

8. Intermolecular Forces (IMF)
Intermolecular forces- describe the interaction between molecules and influence many properties including:
Melting point (solid ↔ liquid)
Boiling point (liquid ↔ gas)
Vapor pressure (gas phase equilibrium of a liquid)

9. IMF & Properties of Water
Water molecules are polar and have strong intermolecular attractions.
Unique properties of water due to IMF:
High surface tension
Capillary action
High boiling point (considering MWt.)
Density of solid < Liquid state

10. Surface Tension  tendency of a liquid to contract in area
 result of molecular attractions
 a sphere has smallest surface area for a given volume
 water has very high surface tension
 explains why needles can float on water

11. Capillary Action  rising of a liquid in a narrow tube
 caused by cohesion and adhesion
 adhesion - attractive force between two unlike substances
 cohesion - attractive force between like substances

12. States of Matter and IMF
IMF defines the energy (temp) requirements for a material to change phase phases.
Some of these relationships can be interpreted by understanding the following graphs:
Vapor Pressure Curves
Phase Diagrams

13. Vapor Pressure:
Equilibrium vapor pressure is the pressure exerted by a vapor at equilibrium with its corresponding liquid at a given temperature (as gas molecules enter the headspace, the Vapor pressure increased)
General trend :  IMF =  Vapor pressure

14. What is Pressure: Useful relationships:
Is defined as force/area.
SI unit = Pascal's
Atmospheric Pressure- collision of air molecules with earth surfaces.  attitude =  atmospheric pressure
Barometer- used to measure atmospheric pressure in mm Hg.
Useful relationships:
1 atm = 760 mmHg = 760 torr= kPa

15. Vapor Pressure Curve:
Describes the boiling point (temp) of a liquid at a given atmospheric pressure.
Vapor pressure increases with increasing temperature.
IMF = Vapor pressure and IMF = Vapor pressure

16. Boiling A liquid boils when the vapor pressure = the external pressure
Normal Boiling point is the temperature a substance boils at 1 atm pressure.
The temperature of a liquid can never rise above it’s boiling point.

17. Changing the Boiling Point – lowering it
Lower the pressure (going up into the mountains).
Lower external pressure requires lower vapor pressure.
Lower vapor pressure means lower boiling point.
Food cooks slower.

18. Changing the Boiling Point – raising it
Raise the external pressure (Use a pressure cooker).
Raises the vapor pressure needed.
Raises the boiling point.
Food cooks faster.

19. Effect of Pressure on Boiling Point
Boiling Point of Water at Various Locations
Location
Feet above sea level
Patm (kPa)
Boiling Point (C)
Top of Mt. Everest, Tibet
29,028 32 70
Top of Mt. Denali, Alaska
20,320
45.3 79
Top of Mt. Whitney, California
14,494
57.3 85
Leadville, Colorado
10,150 68 89
Top of Mt. Washington, N.H.
6,293
78.6 93
Boulder, Colorado
5,430
81.3 94
Madison, Wisconsin
900
97.3 99
New York City, New York 10
101.3
100
Death Valley, California
-282
102.6
100.3

20. Phase Changes Solid Gas Liquid Melting Vaporization Freezing
Condensation

21. Solid Gas Liquid Sublimation Melting Vaporization Freezing
deposition
Freezing
Condensation

22. 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).

23. Phase changes by Name

24. Water