1. Unit 11 – States of Matter & Solutions Chapter 13.2-13.4 & Chapter 15 Unit Test: March 13th

2. Kinetic Molecular Theory All particles are in motion All particles are in motion Kinetic energy increases with temperature Kinetic energy increases with temperature –Particles move faster and farther apart Applies to all states of matter - solids, liquids, and gases Applies to all states of matter - solids, liquids, and gases Ex: concrete expanding on highways Ex: concrete expanding on highways

3. States of Matter - Intermolecular Forces Intermolecular forces act between stable molecules (not like bonding where forces act within molecules) Intermolecular forces act between stable molecules (not like bonding where forces act within molecules) Weak intermolecular forces Weak intermolecular forces –Low boiling points –Most likely in gaseous state Strong intermolecular forces Strong intermolecular forces –High boiling points –Most likely in solid state

4. States of Matter - Solids Definite shape and volume Definite shape and volume Particles vibrate around fixed points Particles vibrate around fixed points Crystalline or amorphous structure Crystalline or amorphous structure Crystalline Structure: Particles arranged in an orderly, repeating 3D pattern Amorphous Structure: Forms when molten material cools too quickly for crystals to form; no definite repeating pattern

5. States of Matter - Liquids Definite volume, but no definite shape Definite volume, but no definite shape Particles vibrate around moving points Particles vibrate around moving points Liquids (& gases) are fluids, meaning they have… Liquids (& gases) are fluids, meaning they have… –Fluidity - ability to flow –Viscosity - resistance to flow –Ex: honey vs. water

6. States of Matter - Liquids Vapor Pressure – pressure exerted by vapor molecules above a liquid when dynamic equilibrium is reached Vapor Pressure – pressure exerted by vapor molecules above a liquid when dynamic equilibrium is reached Dynamic equilibrium – 2 opposite processes occurring at the same rate Dynamic equilibrium – 2 opposite processes occurring at the same rate –Ex: evaporation and condensation Vapor pressure depends on… Vapor pressure depends on… –Temperature ( –Temperature (  T  V.P.) –Strength of IMF ( –Strength of IMF (  IMF  V.P.) Volatile - how easily a fluid evaporates (high vapor pressure) Volatile - how easily a fluid evaporates (high vapor pressure) Vapor WaterAlcohol Vapor Compare the vapor pressure of water and alcohol. What do the arrows mean in these diagrams?

7. States of Matter - Liquids Vapor pressure graphs show vapor pressure vs. temperature Vapor pressure graphs show vapor pressure vs. temperature Standard pressure is 1 atm or 101.3 kPa – indicated by the dark, horizontal line Standard pressure is 1 atm or 101.3 kPa – indicated by the dark, horizontal line –Intersection of curved lines and standard pressure is the normal boiling point

8. States of Matter – Phase Changes Melting/Freezing Melting/Freezing changes between solid and liquid phases Vaporization/Condensation Vaporization/Condensation changes between liquid and gas phases Sublimation/Deposition Sublimation/Deposition changes directly between solid and gas phases

9. States of Matter – Phase Changes Endothermic phase changes absorb energy Endothermic phase changes absorb energy –Sublimation –Vaporization –Melting Exothermic phase changes release energy Exothermic phase changes release energy –Deposition –Condensation –Freezing

10. States of Matter – Phase Changes Vaporization includes 2 processes: Evaporation Evaporation –Only occurs at the surface of the liquid at room temperature –Ex: Sweating cools the human body Boiling Boiling –Occurs throughout the liquid –Boiling point occurs when vapor pressure = atmospheric pressure –Ex: Boiling water Higher altitudes have lower atmospheric pressure, so lower boiling point Higher altitudes have lower atmospheric pressure, so lower boiling point

11. States of Matter - Phase Changes During phase changes, there is no change in kinetic energy - only potential energy increases!!! During phase changes, there is no change in kinetic energy - only potential energy increases!!! –Liquid H 2 O at 0 o C has more kinetic energy than solid H 2 O at 0 o C –Gas H 2 O at 100 o C has more kinetic energy than liquid H 2 O at 100 o C Coke on ice - shows 3 phases of matter

12. States of Matter – Phase Changes Sublimation point Freezing point Boiling point

13. States of Matter – Phase Changes Standard Pressure Line (1 atm) Standard Pressure Line (1 atm) Normal Freezing Point Normal Freezing Point –Temp at which standard pressure meets the solid-liquid curve Normal Boiling Point Normal Boiling Point –Temp at which standard pressure meets the liquid- vapor curve Triple Point Triple Point –Pressure and temperature where all three phases coexist Critical Point Critical Point –Pressure and temperature at which a liquid can no longer exist Normal Boiling PointNormal Freezing Point STP

14. Heterogeneous Mixtures Heterogeneous Mixture: two or more substances physically combined; not uniform throughout Heterogeneous Mixture: two or more substances physically combined; not uniform throughout Ex: Granite, chex mix Ex: Granite, chex mix

15. Unique Properties of Water Hydrogen bonding between water molecules causes: Hydrogen bonding between water molecules causes: –High boiling and melting points –Ability to hold a large amount of heat –High surface tension –Spherical droplets Water is… Water is… –Polar molecule –Universal solvent –Vital for human life Hydration Hydration Transportation of waste Transportation of waste to the kidneys

16. Solutions Homogeneous Mixture: two or more substances physically combined; uniform throughout Homogeneous Mixture: two or more substances physically combined; uniform throughout Composed of… Composed of… –Solute: substance being dissolved (smaller amount) –Solvent: substance that does the dissolving (larger amount) Examples: carbonated drinks, kool-aid, iced tea, ocean Examples: carbonated drinks, kool-aid, iced tea, ocean

17. Solutions - Rate Rate of Solution – how fast a solute will dissolve in a solvent Rate of Solution – how fast a solute will dissolve in a solvent Increase rate of solution by: Increase rate of solution by: –Heating –Stirring –Crushing

18. Solutions - Solubility Solubility – the amount of solute that will dissolve in a solvent at a certain temperature and pressure Solubility – the amount of solute that will dissolve in a solvent at a certain temperature and pressure Soluble: like dissolves like Soluble: like dissolves like –Polar solutes dissolve in polar solvents Ex: water, alcohol, sugar, salt Ex: water, alcohol, sugar, salt –Non-polar solutes dissolve in non-polar solvents Ex: oils, gasoline, diatomic molecules Ex: oils, gasoline, diatomic molecules Insoluble: do not dissolve in each other Insoluble: do not dissolve in each other –Polar and non-polar Ex: oil and water Ex: oil and water

19. Solutions - Solubility Polar vs. Non-polar Molecules: Polar vs. Non-polar Molecules: Non-polar molecules have charges that are evenly distributed, due to shape (ex: any diatomic molecule, gasoline) Non-polar molecules have charges that are evenly distributed, due to shape (ex: any diatomic molecule, gasoline) Polar molecules have a partial positive and a partial negative charge (ex: water) Polar molecules have a partial positive and a partial negative charge (ex: water)

20. Solutions - Solubility Why does “like dissolve like”? Why does “like dissolve like”? –Polar molecules dissolve in other polar molecules because the opposing charges attract and pull the molecule apart –Ex: Dissolving salt (NaCl) in water (H 2 O)

21. Solutions - Solubility Factors that affect solubility Factors that affect solubility –Temperature Solids: ↑ temp ↑ solubility Solids: ↑ temp ↑ solubility Gases: ↑ temp ↓ solubilty Gases: ↑ temp ↓ solubilty –Ex: Fish “breathe” dissolved oxygen gas in the water. When thermal pollution occurs, the increased water temperature makes gases less able to dissolve, and the fish die. –Pressure Solids: no effect Solids: no effect Gases: ↑ pressure ↑ solubility Gases: ↑ pressure ↑ solubility –Ex: Carbonated drinks have dissolved CO 2 because they are under high pressure. When the can is opened, the pressure drops and the gases begin to escape (bubbles).

22. Solutions - Types Unsaturated – solvent contains less solute than it can hold (ex: coffee and sugar) Unsaturated – solvent contains less solute than it can hold (ex: coffee and sugar) Saturated – solvent contains the maximum amount of solute (ex: coffee and sugar) Saturated – solvent contains the maximum amount of solute (ex: coffee and sugar) –If more solute is added, it does not dissolve How could you tell if your coffee was unsaturated or saturated?

23. Solutions - Types Supersaturated – contains more solute than the solvent can normally hold (ex: sodium acetate) Supersaturated – contains more solute than the solvent can normally hold (ex: sodium acetate) –Made by heating the solution to dissolve the excess, then cooling to a lower temperature

24. Solutions – Solubility Graph Solubility graphs show amount of solute vs. temperature Solubility graphs show amount of solute vs. temperature –Above the curve would be supersaturated –The curve itself is at saturation –Below the curve indicates unsaturated solutions –Negative (downward) slope means it is a gas

25. Solutions - Colligative Properties Colligative properties (physical property) depend on the concentration of the particles in the solution Colligative properties (physical property) depend on the concentration of the particles in the solution –Examples: Boiling point elevation: adding salt to water for cooking Boiling point elevation: adding salt to water for cooking Freezing point depression: salting the roads before a freeze, antifreeze in cars, and making homemade ice cream Freezing point depression: salting the roads before a freeze, antifreeze in cars, and making homemade ice cream Osmotic pressure: responsible for plant’s cell wall, sturdiness Osmotic pressure: responsible for plant’s cell wall, sturdiness

26. Solutions - Colligative Properties How does adding a solute change physical properties? How does adding a solute change physical properties? –Solute particles get in the way of the solvent molecules –Makes it harder for the solvent molecules to boil (more energy needed – higher temperature) –Makes it harder for the solvent molecules to freeze (need to release more energy – lower temperature)

27. Solutions - Concentration Concentration is the amount of solute dissolved in a given amount of solvent Concentration is the amount of solute dissolved in a given amount of solvent Described qualitatively as… Described qualitatively as… –Dilute – solution containing a small amount of solute –Concentrated – solution containing a large amount of solute

28. Solutions - Concentration Molarity describes concentration quantitatively Molarity describes concentration quantitatively –The number of moles of solute dissolved in 1 liter of a solution –Ex: molarity of IV fluids is calculated before it is administered to the patient –Equation: Molarity = moles of solute liters of solution liters of solution M = mol L also written as…

29. Solutions - Concentration 1. What is the molarity of 2 mol NaCl in 5 L of solution? 2. How many moles of KBr would be present in 1 L of a 3M solution? 3. What is the volume of a 1.5M solution of HCl that contains 10 moles of HCl? 4. What is the molarity of 2000 mL water containing 49 g H 3 PO 4 ?

30. Solutions - Concentration Dilution adds more solvent to a solution; spreads the solute through a larger volume Dilution adds more solvent to a solution; spreads the solute through a larger volume Example: diluting tea Example: diluting tea Equation is M 1 V 1 = M 2 V 2 where… Equation is M 1 V 1 = M 2 V 2 where… M 1 = initial molarity M 2 = final molarity V 1 = initial volume V 2 = final volume

31. Solutions - Concentration 1. 100 mL of a 3 M solution of NaOH is diluted to 375 mL. What is the new molarity? 2. What volume of 12 M HCl would be required to make 35 mL of 0.5 M HCl solution?