I. Kinetic Molecular Theory KMT

Assumptions of KMT All matter is composed of tiny particles These particles are in constant, random motion. Some particles are moving fast, some are moving slowly. Temperature is a measure of the average Kinetic Energy and is proportional to the average speed of the molecules.

KMT Model http://preparatorychemistry.com/Bishop_KMT_frames.htm Click on the link above to see how particles of matter behave according to the KMT.

II. Intermolecular Forces Liquids & Solids II. Intermolecular Forces

Definition of IMF Attractive forces between molecules. Much weaker than chemical bonds within molecules. a.k.a. van der Waals forces

Types of IMF London Dispersion Forces View animation online.

Types of IMF Dipole-Dipole Forces + - View animation online.

Types of IMF Hydrogen Bonding

Types of IMF

III. Physical Properties Liquids & Solids III. Physical Properties

Liquids vs. Solids IMF Strength Fluid Density Compressible Diffusion Stronger than in gases Y high N slower than in gases SOLIDS Very strong N high extremely slow

Liquid Properties Surface Tension attractive force between particles in a liquid that minimizes surface area

Liquid Properties Capillary Action attractive force between the surface of a liquid and the surface of a solid water mercury

Types of Solids Crystalline - repeating geometric pattern covalent network metallic ionic covalent molecular Amorphous - no geometric pattern decreasing m.p.

Types of Solids Ionic (NaCl) Metallic

Types of Solids Covalent Molecular Covalent Network Amorphous (H2O) (SiO2 - quartz) Amorphous (SiO2 - glass)

Liquids & Solids IV. Changes of State

Phase Changes Most substances can exist in 3 states: Solid Liquid Gas Depends on temperature and pressure

Phase Changes

Phase Changes Each state is referred to as a “Phase” Ice water is a heterogeneous mixture of 2 phases When energy is added or removed, one phase can change into another

Phase Changes Requiring Energy Melting Vaporization Sublimation

Melting Amount of energy needed to melt a substance depends on forces keeping particles together. Melting water requires a high amount of energy because of hydrogen bonding Adding energy allows molecules to move faster, breaking the hydrogen bonds Melting point - Temperature at which the forces holding crystal lattice together are broken and substance becomes liquid

Phase Changes IMF m.p. Melting Point equal to freezing point Which has a higher m.p.? polar or nonpolar? covalent or ionic? polar ionic

Vaporization In liquid water, some particles will have more kinetic energy than others. When the particles have enough energy to overcome the forces of attraction they will escape the liquid as a gas. Vapor - A substance that is liquid at room temperature and becomes gas. Vaporization - the process of changing a liquid to a gas

Vaporization Evaporation - Vaporization that occurs at the surface of a liquid, molecules at the surface gain enough energy to overcome IMF Evaporation is gradual Even at cold temperatures, some molecules have enough energy to break the attractions and become gas. Evaporation is how your body cools itself Water in sweat absorbs heat from your body Water evaporates leaving less heat in your body and a lower ‘average kinetic energy’ (lower temperature)

Phase Changes temp v.p. IMF v.p. Vapor Pressure pressure of vapor above a liquid at equilibrium v.p. depends on temp & IMF temp temp v.p. IMF v.p.

Phase Changes Patm b.p. IMF b.p. Boiling Point Temperature at which the vapor pressure of a liquid equals the external atmospheric pressure depends on Patm & IMF Normal B.P. - b.p. at 1 atm Patm b.p. IMF b.p.

Sublimation Sublimation - when a substance goes from solid directly to gas without becoming a liquid Solid iodine Frozen carbon dioxide (dry ice) Moth balls (p-dichlorobenzene) Solid air fresheners Ice cubes left in a freezer for a long time

Phase Changes the Release Energy Condensation Deposition Freezing

Condensation Condensation – process of changing from gas to liquid When molecule s lose energy, Velocity of the molecules is reduced Intermolecular forces take over Hydrogen bonds form, energy is released (heat comes out) There are different causes for condensation, however, all involve loss of energy: Contact with cold item

Phase Changes Equilibrium trapped molecules reach a balance between evaporation & condensation

Deposition Deposition – changing from gas directly to solid Snowflakes

Freezing Freezing – process of changing from liquid to solid Remove energy from liquid Molecules don’t move past each other any longer Molecules stay in fixed, set position Become solid

Phase Changes Energy Required Energy Released Solid  Liquid Melting or fusion Liquid  Solid freezing Liquid  Gas Vaporization, evaporation or boiling Gas  Liquid condensation Solid  Gas sublimation Gas  Solid solidification

Phase Diagrams Phase diagram shows phase of matter at different temperatures and pressures Each substance is unique X-axis usually temperature Y-axis usually pressure There is usually a “Triple Point” where all three phases can coexist “Critical Point” – temperature and pressure above which substance cannot exist as liquid

Phase Diagram - Water

Phase Diagram CO2

Heating Curves Kinetic Energy Changes – Heat Energy speeds up the molecules. Potential Energy Changes – Heat energy separates the molecules from solid to liquid, liquid to gas.

Heating Curves Temperature Change change in KE (molecular motion) depends on heat capacity Heat Capacity energy required to raise the temp of 1 gram of a substance by 1°C

Phase Change The energy involved in a phase change is calculated using: Heat of Fusion (Hfus) Heat of Vaporization (Hvap)

Heat of Fusion Heat of Fusion is the energy required to change 1 gram of a substance from the solid to the liquid state without changing its temperature. Heat of Fusion is used for calculations involving the phase changes of solid  liquid or liquid solid

Heat of Vaporization Heat of Vaporization (Hvap) energy required to boil 1 gram of a substance at its b.p. EX: sweating, steam burns, the drinking bird