1. Chapter 10 Kinetic Theory (Kinetikos - “Moving”)
Based on the idea that particles of matter are always in motion
The motion has consequences
Behavior of Gases
Physical Properties of Gases
Ideal Gas – an imaginary gas that conforms perfectly to all assumptions

2. Five Assumptions of the KMT
Gases consist of large number of tiny particles
The Particles are in Constant Motion, moving in straight lines.
The collisions between particles & w/ the container wall are elastic.
There are no forces of attraction or repulsion between the particles of a gas.
The average K.E. of the particles is directly proportional to the Kelvin Temperature.
KE = ½ mv2

3. Pressure exerted by the column of air in the atmosphere.
Result of the earth’s gravity attracting the air downward.
Barometer – device used to measure the atmospheric pressure on earth.
Manometer – device used to measure the pressure of a gas in an enclosed container.

5. Physical Properties of Gases
Defined as having indefinite shape and volume.
Gases have mass
Expands to occupy any space available.
Easily compressed
Different gases move easily through each other.
Diffusion – spontaneous mixing of 2 gases.
Low mass = High rate
Effusion – gas passes through tiny opening.
Gases exert pressure
Fluidity – ability to flow.
Low density

6. Real Gas
Gas that does not behave completely to the assumption of the KMT.
Deviation from ideal behavior:
High Pressure
Low Temperature
Polar molecules

8. 10-2 Liquids Definite volume but no definite shape.
Liquids are fluids.
Liquids attractions are a result of the IMF’s.

9. Liquid Properties Surface Tension Capillary Action
Force that tends to pull adjacent parts of a liquid’s surface together, decreasing the surface area to the smallest possible size.
Imbalance of forces at the surface of a liquid.
Capillary Action
Attraction of the surface of a liquid to the surface of a solid.

10. Viscosity
“Friction” or Resistance to motion, that exist between molecules in a liquid.
High Viscosity = Low Flow
Stronger IMF = Higher Viscosity
Increase KE = Low Viscosity

11. Evaporation/Boiling Vaporization Evaporation Boiling
Process in which a liquid changes to a gas.
Evaporation
Process in which particles escape the surface of a non-boiling liquid and enter the gas phase.
This is caused by a greater KE at the surface of the liquid.
Boiling
Conversion of a liquid to a gas within the liquid as well as at its surface.

12. 10-3 Solids
Definite shape and definite volume.

13. Crystals Crystals have an ordered, repeated structure.
The smallest repeating unit in a crystal is a unit cell.
Three-dimensional stacking of unit cells is the crystal lattice.
Amorphous Solids
Lack internal order but yet exhibit a solid like substance.
Jello is similar but it’s considered a colloidal suspension.

15. Structures of Solids
Unit Cells

16. Structures of Solids
Unit Cells

17. Structures of Solids
Close Packing of Spheres

18. Close Packing of Spheres
A crystal is built up by placing close packed layers of spheres on top of each other.
There is only one place for the second layer of spheres.
There are two choices for the third layer of spheres:
Third layer eclipses the first (ABAB arrangement). This is called hexagonal close packing (hcp).
Third layer is in a different position relative to the first (ABCABC arrangement).

19. Close Packing of Spheres
Each sphere is surrounded by 12 other spheres (6 in one plane, 3 above and 3 below).
Coordination number: the number of spheres directly surrounding a central sphere.

20. Crystal Bonding Metallic Solids (mobile valence electrons)
Low to High melting points
Metallic bonds hold the particles together
Molecular Solids (lowest melting pts)
Low melting points
Intermolecular forces hold the particles together
Ionic Solids (hard, brittle and non-conducting)
High melting points
Strong electrostatic force of attraction
Covalent – Network Solids (strong covalent bonds between neighboring atoms)
Atoms covalently bonded to the same type of atoms

21. 10-4 Changes of State Phase change
Conversion of a substance from one of the 3 physical states of matter to the other.
Always involves a change in energy.

22. Equilibrium Equilibrium ()
Dynamic condition in which 2 opposing changes occur at equal rates in a closed system.
Components under equilibrium
Phase – any part of the system that has uniform composition and properties.
System – sample of matter being studied.
Concentration - #particles per unit of volume

23. Phase Change Evaporation/Condensation Freezing/Melting
Evaporation – rate in which a liquid changes to a gas under its boiling point.
Condensation – rate in which a gas changes to a liquid.
Phase change : Evaporation  Condensation
Liquid + Heat  Vapor
Vapor  Liquid + Heat
Freezing/Melting
Freezing – rate in which a liquid changes to a solid.
Melting – rate in which a solid changes to a liquid.
Phase change : Freezing  Melting
Solid + Heat  Liquid
Liquid Solid + Heat

24. Sublimation/Deposition
Phase change that occurs when a solid changes to a gas without passing through the liquid phase.

25. Possible Changes of State
Name
Example
Gas  Liquid
Condensation
H2O(g)  H2O(l)
Liquid  Gas
Vaporization
Br(l)  Br(g)
Liquid  Solid
Freezing
H2O(l)  H2O(s)
Solid  Liquid
Melting
H2O(s)  H2O(l)
Solid  Gas
Sublimation
CO2(s)  CO2(g)

26. Boiling
Conversion of a liquid to a vapor, when the vapor pressure of the liquid is equal to the atmospheric pressure.
Vapor Pressure – Amount of pressure caused by the vapor of a liquid in a closed container.
Boiling Point – Temperature at which a liquid’s vapor pressure equals the atmospheric pressure.
Normal Boiling Point – Temperature at which a liquid boils at Standard Pressure.

28. Factors Affecting Boiling
2 Factors that cause boiling:
Lowering the atmospheric pressure, by placing the liquid in a vacuum.
Increasing the vapor pressure, by increasing the temperature of the liquid.

29. Phase Diagrams
Graph of Temperature vs. Pressure that indicates points in which a substance will be a gas, liquid or a solid.
Triple Point – Temperature and Pressure at which a substance has all three phases at equilibrium.
Critical Point – Point in which a substance can’t exist in the liquid state.

30. Phase Graph of Water

31. Phase Diagram

32. Molar Heats (Enthalpy)
Molar Heat of Fusion/Solidification
Amount of heat needed to change 1 mole of a substance from a liquid to a solid or solid to a liquid.
Solid  Liquid (Molar Heat of Fusion)
Liquid  Solid (Molar Heat of Solid)
Water: Molar Heat of Fusion
(Hfus) = 6.01 kJ/mol

33. Molar Heats Cont. Molar Heat of Condensation/Vaporization
Amount of heat needed to change 1 mole of a substance from a liquid to a gas or a gas to a liquid.
Gas  Liquid (Molar Heat of Condensation)
Liquid  Gas (Molar Heat of Vaporization)
Water: Molar Heat of Vaporization
(Hvap) = 40.7 kJ/mol

34. Molar Heat Problem
Determine the amount of heat needed to melt 100g of ice at 0oC.
Determine the amount of heat needed to change 100g of liquid water to steam.

35. 10-5 Water Water is present in a large abundance throughout our life.
70%-75% earth’s surface is water
60%-90% of the mass of most living things is water.

36. Water’s the Exception Water expands when it freezes
Less dense than water
Reason for ice floating
3.98oC water begins to expand due to crystal formation in water.

37. Water’s Hydrogen Bonding
Two Types of Strong Interactions
Cohesive forces – between molecules of the same type
Adhesive forces – between different types

38. Adhesive & Cohesive
Water in a tube exhibits a curved surface called a meniscus.
Adhesive forces occur between the glass and water
Drawing the water up along the glass
Cohesive forces in the water help hold the curve in the water level