2. Chapter 14: Solutions and Their Properties If you’re not part of the solution, you’re part of the precipitate!

3. Properties of Solutions

4. Physical Properties of Solutions Have you ever wondered... Why antifreeze keeps water from freezing? Why salt causes ice to melt? Why cooks add salt to boiling water? Why root beer foams only when poured? What force opposes gravity to allow water to climb up a tree?

5. Solutions Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent.

6. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. nonelectrolyte weak electrolyte strong electrolyte

7. “like dissolves like” Two substances with similar intermolecular forces are likely to be soluble in each other. non-polar molecules are soluble in non-polar solvents CCl 4 in C 6 H 6 polar molecules are soluble in polar solvents C 2 H 5 OH in H 2 O ionic compounds are more soluble in polar solvents NaCl in H 2 O or NH 3 (l)

8. Energy Changes in Solution To determine the enthalpy change, we divide the process into 3 steps. 1.Separation of solute particles. 2.Separation of solvent particles to make ‘holes’. 3.Formation of new interactions between solute and solvent.

9. Three types of interactions in the solution process: solvent-solvent interaction solute-solute interaction solvent-solute interaction  H soln =  H 1 +  H 2 +  H 3

10. Enthalpy Changes in Solution The enthalpy change of the overall process depends on  H for each of these steps. Start End Start

11. Calculating  H solution (enthalpy of solution) KF (s) -> K + (g) + F-(g)  E = +821 kJ K + (g) + F - (g) + H 2 O -> K + (aq) + F - (aq)  E=-819 kJ So Net - KF (s) -> K + (aq) + F - (aq)  H solution = +2kJ

12. Enthalpy changes during dissolution The enthalpy of solution,  H soln, can be either positive or negative.  H soln =  H 1 +  H 2 +  H 3  H soln (MgSO 4 )= -91.2 kJ/mol --> exothermic  H soln (NH 4 NO 3 )= 26.4 kJ/mol --> endothermic

13. Why do endothermic processes sometimes occur spontaneously? Some processes, like the dissolution of NH 4 NO 3 in water, are spontaneous at room temperature even though heat is absorbed, not released.

14. Enthalpy Is Only Part of the Picture Entropy is a measure of: Dispersal of energy in the system. Amount of disorder in system The solution has greater entropy or disorder, and  is the favored state

15. Entropy changes during dissolution Each step also involves a change in entropy. 1.Separation of solute particles. 2.Separation of solvent particles to make ‘holes’. 3.Formation of new interactions between solute and solvent.

16. Degree of saturation Saturated solution  Solvent holds as much solute as is possible at that temperature.  Undissolved solid remains in flask.  Dissolved solute is in dynamic equilibrium with solid solute particles.

17. Degree of saturation Supersaturated  Solvent holds more solute than is normally possible at that temperature.  These solutions are unstable; crystallization can often be stimulated by adding a “seed crystal” or scratching the side of the flask.

18. Gases in Solution In general, the solubility of gases in water increases with increasing mass. Why? Larger molecules have stronger dispersion forces.

19. Gases in Solution The solubility of liquids and solids does not change appreciably with pressure. But, the solubility of a gas in a liquid is directly proportional to its pressure. Increasing pressure above solution forces more gas to dissolve.

20. Henry’s Law What happens to the solubility of carbon dioxide in a bottle of soda when the pressure is reduced?

21. Pressure and Solubility of Gases The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution (Henry’s law). c = kP c is the concentration (M) of the dissolved gas P is the pressure of the gas over the solution k is a constant (mol/Latm) that depends only on temperature low P low c high P high c

22. Chemistry In Action: The Killer Lake Lake Nyos, West Africa 8/21/86 CO 2 Cloud Released 1700 Casualties Trigger? earthquake landslide strong Winds

23. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature. Solubility is measured as the mass of solute dissolved in 100 g of solvent at a given temperature

24. Temperature The opposite is true of gases. Higher temperature drives gases out of solution.  Carbonated soft drinks are more “bubbly” if stored in the refrigerator.  Warm lakes have less O 2 dissolved in them than cool lakes.

25. Mass Percentage Mass % of A = mass of A in solution total mass of solution  100

26. moles of A total moles in solution X A = Mole Fraction (X) In some applications, one needs the mole fraction of solvent, not solute—make sure you find the quantity you need!

27. mol of solute L of solution M = Molarity (M) Because volume is temperature dependent, molarity can change with temperature.

28. mol of solute kg of solvent m = Molality (m) Because neither moles nor mass change with temperature, molality (unlike molarity) is not temperature dependent.

29. Mass/Mass Moles/Moles Moles/Mass Moles/L

30. Solution Concentration Complete the following table: Compound / MW MolalityWeight %Mole Fraction KI / 116 g/mol 0.15 C 2 H 5 OH / 46.07g/mol 3.0%

31. What is the molality of a 5.86 M ethanol (C 2 H 5 OH) solution whose density is 0.927 g/mL? m =m = moles of solute mass of solvent (kg) M = moles of solute liters of solution Assume 1 L of solution: 5.86 moles ethanol = 270 g ethanol 927 g of solution (1000 mL x 0.927 g/mL) mass of solvent = mass of solution – mass of solute m =m = moles of solute mass of solvent (kg) = 5.86 moles C 2 H 5 OH 0.657 kg solvent = 8.92 m mass of solvent = mass of solution – mass of solute = 927 g – 270 g = 657 g = 0.657 kg

32. Changing Molarity to Molality If we know the density of the solution, we can calculate the molality from the molarity, and vice versa.

33. Convert % mass to Molarity What is the Molarity of a 95% acetic acid solution? (density = 1.049 g/mL) If you assume 1 L, that amount of solution = 1049 g 95% of the solution is acetic acid 1049 g solution x 0.95 = 997 g solute 997 g X 1 mol/60.05 g = 16.6 mol solute Since we assumed 1 L, that’s 16.6 mol / 1 L or 16.6 M

34. Colligative Properties Colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. Among colligative properties are  Vapor pressure lowering  Boiling point elevation  Melting point depression  Osmotic pressure

35. Vapor Pressures of Pure Water and a Water Solution The vapor pressure of water over pure water is greater than the vapor pressure of water over an aqueous solution containing a nonvolatile solute. Solute particles take up surface area and lower the vapor pressure

36. Vapor Pressure As solute molecules are added to a solution, the solvent becomes less volatile (has decreased vapor pressure). Solute-solvent interactions contribute to this effect.

37. Vapor Pressure Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

38. Lowering Vapor Pressure Raoult’s Law: Where: P A = vapor pressure with solute, P A  = vapor pressure without solute (pure solvent), and  A = mole fraction of A (the pure solvent). Colligative Properties

39. Boiling Point Elevation and Freezing Point Depression Solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

40. Boiling Point Elevation The change in boiling point is proportional to the molality of the solution:  T b = K b  m where K b is the molal boiling point elevation constant, a property of the solvent.  T b is added to the normal boiling point of the solvent.

41. Freezing Point Depression The change in freezing point can be found similarly:  T f = K f  m Here K f is the molal freezing point depression constant of the solvent.  T f is subtracted from the normal freezing point of the solvent.

42. Boiling-Point Elevation Molal boiling-point-elevation constant, K b, expresses how much  T b changes with molality, m S : Decrease in freezing point (  T f ) is directly proportional to molality (K f is the molal freezing-point-depression constant): Colligative Properties In both equations,  T does not depend on what the solute is, but only on how many particles are dissolved.

43. Colligative Properties of Electrolytes Because these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) show greater changes than those of nonelectrolytes. e.g. NaCl dissociates to form 2 ion particles; its limiting van’t Hoff factor is 2.

44. Colligative Properties of Electrolytes However, a 1 M solution of NaCl does not show twice the change in freezing point that a 1 M solution of methanol does. It doesn’t act like there are really 2 particles.

45. van’t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions.

46. van’t Hoff Factor Some Na + and Cl − reassociate as hydrated ion pairs, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

47. The van’t Hoff Factor Reassociation is more likely at higher concentration. Therefore, the number of particles present is concentration dependent.

48. The van’t Hoff Factor We modify the previous equations by multiplying by the van’t Hoff factor, i  T f = K f  m  i i = 1 for non-elecrtolytes

49. Osmosis Semipermeable membranes allow some particles to pass through while blocking others. In biological systems, most semipermeable membranes (such as cell walls) allow water to pass through, but block solutes.

50. Osmosis In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration). Water tries to equalize the concentration on both sides until pressure is too high.

51. Osmosis Osmotic pressure, , is the pressure required to stop osmosis: Colligative Properties

52. Molar Mass from Colligative Properties We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound.

53. Osmosis in Blood Cells If the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic. Water will flow out of the cell, and crenation results.

54. Osmosis in Cells If the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic. Water will flow into the cell, and hemolysis results.

55. Properties of Solutions