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Chapter 15

Chapter 15. Solutions. Solutions. Have things dissolved? Solutions involve all states of matter Gas and gas (Air) Gas and liquid (Carbonated drinks) Liquid and solid (Ocean) Liquid and liquid (Alcohol and water) Solid and solid (Alloy) The solute is dissolved into the solvent.

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Chapter 15

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  1. Chapter 15 Solutions

  2. Solutions • Have things dissolved? • Solutions involve all states of matter • Gas and gas (Air) • Gas and liquid (Carbonated drinks) • Liquid and solid (Ocean) • Liquid and liquid (Alcohol and water) • Solid and solid (Alloy) • The solute is dissolved into the solvent. • The solute is equally distributed in the solvent. • Homogeneous (same throughout)

  3. How Solutes Dissolve • Solvation is the process of forming a solution. • How particles dissolve • Many ionic compounds come apart to form ions • Dissociation • Ex. NaCl becomes Na+ and Cl- in water • Molecules do not come apart • Ex. Sucrose in water • The sucrose sugar cube will break apart but the actual molecules will not. • The water attaches to the sucrose molecule through hydrogen bonding.

  4. Dissolving Liquids in Liquids • Miscible • Two liquids that are soluble in each other • Alcohol and water • Immiscible • Two liquid that are not soluble in each other • Oil and Water • Why does this occur? • Like dissolves like (Polar dissolves Polar)

  5. Increasing the Rate of Solvation • Heat • Stirring • Increase surface area

  6. Solubility • Refers to the maximum amount of solute that will dissolve in a solvent. • Types of solution are based on solubility • Saturated solution • Contains the maximum amount of solute for a given amount of solvent. • Unsaturated solution • Contains less than the maximum amount of solute for a given amount of solvent. • Supersaturated Solution • Contains more than the maximum amount of solute for a given amount of solvent. • Achieved by increasing temperature • If disturbed the excess solute will crystalize

  7. Henry’s Law • States that the solubility of a gas in a liquid is proportional to the pressure of the gas above the liquid. • Ex. Why are coke’s bottled under pressure? S1 S2 S is the solubility (g/L) P1 P2 P is the pressure

  8. Examples for Henry’s Law • Pg 461 #1 • S1 = 0.55 g / 1 L S2 = ? • P1 = 20.0 kPa P2 = 110.0 kPa • 0.55 g/L S2 20.0 kPa 110.0 kPa • S2 = (0.55 g/L)(110.0 kPa) 20.0 kPa • S2 = 3.0 g/L

  9. Concentration • The amount of solute dissolved in a given amount of solvent. • % by mass • Pg. 463 • % mass mass of solute x 100% mass of solution

  10. % by Mass Examples • Pg 463 #8 • % NaHCO3 20 g NaHCO3x 100% 600 g H2O + 20 g NaHCO3 % by mass of NaHCO3 = 3% • Percent by volume works the same way

  11. Molarity mol solute M L solution • Must be in liters • Example • Pg 465 # 14 40.0 g C6H12O60.148 M 1.5 L

  12. Preparing Molar Solutions • Pg. 466 #17 0.10 mol CaCl2x 110.984 g CaCl211 g CaCl2 1 L 1 mol CaCl2 • Molarity is mol per 1 L. In this problem we have 0.10 mol CaCl2 per 1 L. Next, you must convert moles to grams and we use molar mass.

  13. Homework • Pg. 484 # 64,65,69, 70, 71, 76, 77, 78

  14. Dilution • Adding water to a known solution • M1V1 = M2V2 • M is molarity • V is volume

  15. Dilution Examples • Pg. 468 #21 • M1 = 3.00 M KI M2 = 1.25 M KI • V1 = ? V2 = 0.300 L • V1 = M2V2 / M1 • V1 = .125 L

  16. Molality • mmol solute 1 kg solvent • Example: Pg 469 #24 10.0gNa2SO4x 1 mol Na2SO40.0704mNa2SO4 142.05 g Na2SO4

  17. Mole Fractions • XAmol A mol A + mol B

  18. Example: 25.5g NaBr in 150.0gH2O. What is the mole fraction of NaBr? 25.5gNaBr x 1 molNaBr .2478 molNaBr 102.89gNaBr 150.0gH2O x 1 mol H2O 8.324 mol H2O 18.02gH2O XNaBr .2478 molNaBr .2478 molNaBr + 8.324 mol H2O XNaBr = 0.00983

  19. Homework • Pg 485 82-84

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