1. Which of the following concentration measures will change in value as the temperature of a solution changes? a) Mass percent b) Mole fraction c) Molality d) Molarity

2. What do you remember? Saturated? Unsaturated? Supersaturated? Solubility? Like Dissolves like? Solvation / Hydration? Has max. amount solute that can dissolved Less than max. amt. solute More than max amt. solute Maximum amount solute that can dissolve in solvent

3. Solvation: n process of solvent particles pulling solute particles into solution and forming a sphere around them. n Hydration: water is the solvent

4. Enthalpy (∆H soln ) of Solution: the overall change in heat when a solute is dissolved in a solvent.

5. Three types of interactions in the solution process: Separating solute particles (endothermic) Overcoming solvent IMF’s to make room for solute solvent-solute interaction (usually exothermic)  H soln =  H 1 +  H 2 +  H 3

6. Factors Affecting Solubility: n temperature n pressure (gas solute) n polarity

7. Henry’s Law: n the amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution: C = k∙P C = solute concentration k = constant P = solute pressure

8. Problem: n 5.0 atm of pressure is over a bottled soft drink. What is the concentration of CO 2 in the unopened bottle? k =.031 M C = k·P = 0.155 M

9. Colligative Property n properties that depend upon the solute concentration, but not its identity.

10. Colligative Properties n vapor pressure n boiling point n freezing point n Osmotic pressure

11. 1 M NaCl n # moles NaCl in 1 L? n # moles of particles (ions) in 1 L? one two

12. 1 M sucrose n # moles sucrose in 1 L? n # moles of particles in 1 L? one

13. Vapor Pressure Reduction n a reduction in a solvent’s vapor pressure from the addition of a nonvolatile solute

14. Vapor Pressures of Solutions n Nonvolatile solutes lower the vapor pressure of solvents. n Raoult’s Law: P soln = X solvent ∙ P° solvent where P soln = observed vapor pressure of solution X solvent = mole fraction of solvent X solvent = mole fraction of solvent P° solvent = vapor pressure of pure solvent P° solvent = vapor pressure of pure solvent

15. Question: n If solvent mole fraction is 0.75, what happens to the vapor pressure compared to the pure solvent? = ¾ of the original value

16. A Solution Obeying Raoult’s Law

17. Nonideal Solutions n n Liquid-liquid solutions where both components are volatile. n n Modified Raoult’s Law: n n Nonideal solutions behave ideally as the mole fractions approach 0 and 1.

18. Vapor Pressure for a Solution of Two Volatile Liquids

19. Summary of the Behavior of Various Types of Solutions Interactive Forces Between Solute (A) and Solvent (B) Particles  H soln  T for Solution Formation Deviation from Raoult’s Law Example A  A, B  B  A  B Zero None (ideal solution) Benzene- toluene A  A, B  B < A  B Negative (exothermic) PositiveNegative Acetone- water A  A, B  B > A  B Positive (endothermic) NegativePositive Ethanol- hexane

20. A salt solution sits in an open beaker. Assuming constant temperature and pressure, what will happen to the vapor pressure of the solution? a) It will increase over time. b) It will decrease over time. c) It will stay the same over time.

21. In a 0.1 molar solution of NaCl in water, which one of the following will be closest to 0.1? a) The mole fraction of NaCl b) The mass fraction of NaCl c) The mass percent of NaCl d) The molality of NaCl e) All of these are about 0.1

22. Concept Check For each of the following solutions, would you expect it to be relatively ideal (with respect to Raoult’s Law), show a positive deviation, or show a negative deviation? a)Hexane (C 6 H 14 ) and chloroform (CHCl 3 ) b)Ethyl alcohol (C 2 H 5 OH) and water c)Hexane (C 6 H 14 ) and octane (C 8 H 18 )

23. Boiling Point Elevation: ΔT b n the boiling point of the solution minus the boiling point of the pure solvent. n ΔT b = K b ·m·i

24. Problem: n What boiling point elevation occurs if 200 g. sucrose (C 12 H 22 O 11 ) is added to 500 g. of water? K b for water = 0.52 °C/m. 0.61 °C

25. Freezing Point Depression: ΔT f n n the freezing point of the pure solvent minus the freezing point of the solution. n n ΔT f = K f ·m·i

26. Problem: n What is the freezing point depression of 0.500 kg ethylene glycol (C 2 H 6 O 2 ) in 0.500 kg H 2 O? K f for water=1.86°C/m. 30 °C

27. Molar Mass problem: n The freezing point of water is −0.39°C when 39 g. of a molecular solute are dissolved in 475 g. of water. What is the solute’s molar mass? K f of water = 1.86°C /m 391 g/mol

28. Osmosis: a flow of solvent through a semipermeable membrane from a less concentrated solution to a more concentrated solution. a flow of solvent through a semipermeable membrane from a less concentrated solution to a more concentrated solution.

29. Osmotic Pressure (Π) n pressure created by the movement of solvent. n Π = (i)nRT/V = i∙M∙R∙T n Π = (i)nRT/V = i∙M∙R∙T Π = osmotic pressure M = molarity R = gas law const. T = temperature (K)

30. Molar Mass problem: n When 33.4 mg of a molecular compound is dissolved in 10.0 mL of water at 25.0°C, the solution has an osmotic pressure of 558 torr. Calculate the molar mass of this compound. 111 g/mol

31. Van’t Hoff factor (i) n The number of particles each salt particle dissociates into. n Examples: – Sugar: i = 1 – – NaCl: i = 2 – Fe 3 (PO 4 ) 2 – Fe 3 (PO 4 ) 2 i = ? 5

32. Ion Pairing n The small percentage of ions that do not dissociate in solution. –Exists to some degree in all solutions –More of a factor as solute concentrations increase

33. As the concentration of a NaCl solution increases, ion pairing is expected to: a) increase. b) decrease. c) stay the same.

34. Which of the following 0.05 m solutions is expected to exhibit the greatest ion pairing? a) NaCl b) MgCl 2 c) FeCl 3

35. Colloid: n a mixture of suspended particles that do not settle out.