CHAPTER 10. States of Matter

Transcription

1 CHAPTER 10 States of Matter

2 Kinetic Molecular Theory Kinetikos - Moving Based on the idea that particles of matter are always in motion The motion has consequences Explains the behavior of Gases, Liquids, and Solids Ideal Gas an imaginary gas that conforms perfectly to all assumptions of the KMT

3 Five Assumptions of the KMT 1. Gases consist of large number of tiny particles that are far apart 2. The collisions between particles & w/ the container wall are elastic (no net loss of KE) 3. The Particles are in constant, rapid, random motion, moving in straight lines. 4. There are no forces of attraction or repulsion between the particles of a gas. 5. The average K.E. of the particles is directly proportional to the Kelvin Temperature. Remember: KE = ½ mv 2

4 Atmospheric Pressure 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 Barometer Manometer

6 Physical Properties of Gases Have indefinite shape and indefinite volume. Has mass Expands to occupy any space available. Easily compressed Gases exert pressure Fluidity ability to flow (just like liquids!) Low density

7 Mixing of Gases Different gases move easily through each other. Diffusion spontaneous mixing of 2 gases. Low mass = High rate Effusion gas passes through tiny opening.

8 Real Gas Gas that does not behave completely to the assumption of the KMT. Reasons gases deviate from ideal behavior: High Pressure Low Temperature Polar molecules like water and ammonia

9 Liquids Definite volume but no definite shape. Liquids are fluids (just like gases!) Molecules are held together as a result of the IMF s.

10 Liquid Properties Surface Tension Force that tends to pull adjacent parts of a liquid s surface together, decreasing the surface area to the smallest possible size. Result is that the surface acts like an elastic film.

11 Capillary Action Water forms a meniscus (curved surface) by using cohesive and adhesive forces. This process is called capillary action Liquids use cohesive and adhesive forces to rise within a cylinder Examples: Oil moving up a lamp wick, water creeping up a towel

12 Capillary Action The thinner the tube, the taller the column of water will be!

13 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

14 Evaporation/Boiling Vaporization 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.

15 Solids Definite shape and definite volume Definite melting point High density Incompressible Low rate of diffusion Still happens, but millions of times slower than liquids

16 Crystals vs. Amorphous Solids 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. AKA supercooled liquids Examples: Glass, plastic

17 Types of Crystals

18 Structures of Solids Unit Cells

19 Structures of Solids Unit Cells

20 Structures of Solids Close Packing of Spheres

21 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).

22 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.

23 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

24 Crystal Bonding Ionic Solids (hard, brittle and nonconducting) High melting points Strong electrostatic force of attraction Covalent Network Solids (strong covalent bonds between neighboring atoms) High melting points Atoms covalently bonded to the same type of atoms

25 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.

26 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

27 Phase Changes Sublimation Melting (Add KE) Vaporization (Add KE) Solid Liquid Gas Freezing Condensation (Remove KE) (Remove KE) Deposition

28 Phase Changes Evaporation/Condensation 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

29 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.

30 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.

31 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 at which a substance can t exist in the liquid state.

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33 Phase Diagram

34 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 ( H fus ) = 6.01 kj/mol

35 Molar Heats (Enthalpy) 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 ( H vap ) = 40.7 kj/mol

36 Molar Heat Problems Determine the amount of heat needed to melt 100g of ice at 0 o C. 100g 1mol 6.01kJ x x =33.4 kj 1 18g 1mol o Determine the amount of heat needed to change 100g of liquid water to steam. 100g 1mol 40.7kJ x x =226 kj 1 18g 1mol

37 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. Because water exhibits hydrogen bonding, it behaves differently than most covalent compounds!

38 Water s the Exception Water expands when it freezes Hydogen bonds force a hexagonal crystal pattern Ice floats (less dense) At 3.98 o C water begins to expand due to crystal formation in water.