1. Daniel L. Reger Scott R. Goode David W. Ball http://academic.cengage.com/chemistry/reger Chapter 12 Solutions

2. There are a number of ways to express concentration. You have seen: molarity; mole fraction; mass percentage: for expressing the composition of a compound; and can be used for solutions, as well. normality Solution Concentration

3. All concentration units are fractions. The numerator contains the quantity of solute. The denominator is the quantity of either solution or solvent. They differ in the units used to express these two quantities. Solution Concentration

4. Units of Concentration Used Earlier

5. Mass Percent Composition

6. A solution is prepared by dissolving 3.00 g of NaCl (molar mass = 58.44 g/mol) in 150 g of water. Express its concentration as mass percent. Example: Percent Composition

7. Molality Molality (m or molal) is defined as

8. Example: Calculate Molality What is the molality of a solution prepared by dissolving 3.00 g NaCl (molar mass = 58.44 g/mol) in 150 g of water?

9. Express the concentration of a 3.00% H 2 O 2 solution as (a) molality; (b) mole fraction. Example: Concentration Conversion

10. Test Your Skill Calculate (a) the molality, and (b) the mole fraction of alcohol (C 2 H 5 OH; molar mass = 46.07 g/mol) in a wine that has an alcohol concentration of 7.50 mass percent.

11. Conversion of most concentration units to molarity usually involve using the density of the solution to convert units of mass to units of volume. The density of a 12.0% sulfuric acid (H 2 SO 4 ; molar mass = 98.08 g/mol) is 1.080 g/mL. What is the molarity of this solution? Example: Conversion to Molarity

12. For most substances, there is a limit to the quantity of solute that dissolves in a specific quantity of solvent. A dynamic equilibrium exists between the solute particles in solution and the undissolved solute. Principles of Solubility

13. Normality Normality, (N) is defined as the number of equivalents per liter of solution. Problem – what are equivalents – depends on reaction Not used very often One place that normality is sometimes encountered is with acid-base chemistry. How many protons are release by an acid defines the equivalents EX: HCl – releases 1 H + so… 1M = 1N H 2 SO 4 – releases 2H + so… 1M = 2N We will not be using normality, but you might encounter it in other chemistry classes!!

14. Solubility is the concentration of solute that exists in equilibrium with an excess of that substance. A saturated solution has a concentration of solute equal to its solubility. Definitions

15. Definitions (continued) An unsaturated solution is one that has a solute concentration less than the solubility. A supersaturated solution is one with a solute concentration that is greater than the solubility. Supersaturation is an unstable condition.

16. Many spontaneous processes are exothermic. Enthalpy of solution: the  H that accompanies the dissolution of one mole of solute. When a solid and a liquid form a solution, the enthalpy change arises mainly from changes in the intermolecular attractions. Solute-Solvent Interactions

17. The Solution Process Steps 1 and 2 are endothermic; step 3 is exothermic.

18. A decrease in enthalpy is an important factor in causing spontaneous change. However, many endothermic processes are spontaneous, suggesting another contribution to spontaneity. In increase in disorder also favors spontaneous change. Spontaneity

19. An example of increasing disorder as a driving force is illustrated by the mixing of gases. Spontaneous Mixing of Gases (a) separated gases(b) spontaneously mixed

20. An increase in disorder generally accompanies the mixing of molecules in the formation of a solution. Ammonium nitrate is very soluble in water because of the increase in disorder upon mixing, even though the process is quite exothermic (  H soln = +26.4 kJ/mol). Disorder and Spontaneity

21. Relative solubilities can often be predicted by comparing the relative strengths of the intermolecular attractions of solute-solute, solvent- solvent, and solute-solvent interactions. Solubility of Molecular Compounds

22. Like Dissolves Like In general, substances that have similar intermolecular forces have strong solute- solvent interactions and tend to form solution.

23. (a) Is iodine (I 2 ) more soluble in water or in hexane (C 6 H 14 )? (b) Is methanol (CH 3 OH) more soluble in water or in octane (C 8 H 18 )? Example: Relative Solubility

24. Hydration is the interaction of water molecules with ions, and is very exothermic. Interaction of Ions with Water

25. When an ionic compound dissolves in water, disorder changes because: separating the ions increases disorder; separating the water molecules increases disorder; hydrating the ions, which restricts some water molecules, decreases disorder. A few examples are known where disorder decreases on dissolving ionic compounds. Ionic Compounds in Water

26. Pressure and Solubility Pressure has very little effect on the solubilities of liquids and solids. The solubility of gases in a liquid depends on the pressure of the gas. Henry’s Law: The solubility of a gas is directly proportional to its partial pressure at any given temperature: C = kP

27. Henry’s Law Constants in Water for Various Gases (molal/atm) Gas0°C0°C20°C40°C60°C CO 2 7.60 x 10 -2 3.91 x 10 -2 2.44 x 10 -2 1.63 x 10 -2 C2H4C2H4 1.14 x 10 -2 5.60 x 10 -3 3.43 x 10 -3 --- He4.22 x 10 -4 3.87 x 10 -4 4.10 x 10 -4 N2N2 1.03 x 10 -3 7.34 x 10 -4 5.55 x 10 -4 4.85 x 10 -4 O2O2 2.21 x 10 -3 1.43 x 10 -3 1.02 x 10 -3 8.71 x 10 -4

28. Water at 20  C is saturated with air that contains CO 2 at a partial pressure of 8.0 torr. What is the molal concentration of CO 2 in the solution? Henry’s Law Calculation

29. Experiments show that the way solubility changes with temperature depends on the sign of the enthalpy of solution. Solubility increases with increasing temperature if  H soln is positive (endothermic). Solubility decreases with increasing temperature if  H soln is negative (exothermic). Solubility and Temperature

30. K 2 Cr 2 O 7 (s) → K 2 Cr 2 O 7 (aq)  H = +66.5 kJ/mol The solubility increases with increasing temperature when  H soln is positive. Solubility and  H soln

31. Temperature Dependence on Solubility

32. Colligative property: Any property of a solution that changes in proportion to the concentration of solute particles. Many colligative properties are directly related to the lowering of solvent vapor pressure by the presence of solute particles. Colligative Properties of Solutions

33. Effect of Solute on Evaporation The rate of evaporation of solvent in a solution is lower than that of the pure solvent. Solute particles block opportunities for solvent particles to enter the vapor phase.

34. Raoult’s Law Raoult’s law: The vapor pressure of solvent above a dilute solution equals the mole fraction of the solvent times the vapor pressure of the pure solvent. P solv =  solv P  solv Another form of this equation gives the lowering of the vapor pressure.  P solv =  solute P  solv

35. At 27  C, the vapor pressure of benzene is 104 torr. What is the vapor pressure of a solution that has 0.100 mol of naphthalene in 9.90 mol of benzene? Example: Raoult’s Law

36. Boiling Point Elevation Because a solute lowers the vapor pressure of the solvent, it raises the boiling point of the solution. Below, the concentration of the solution is increasing from (a) to (e).

37. Boiling Point Elevation The boiling point elevation is  T b = mk b where m is the molal concentration and k b is the boiling point constant for the solvent. SolventB.P. (°C)k b (°C/m) Acetic acid117.903.07 Benzene80.102.53 Water100.00.512

38. Freezing Point Depression Solute particles interfere with the ability of solvent particles to form a crystal and freeze. Thus, it takes a lower temperature to freeze solvent from a solution than from the pure solvent. This is freezing point depression.

39. Freezing Point Depression The freezing point depression is  T f = mk f where m is the molal concentration and k f is the freezing point constant for the solvent. Solvent Freezing Pt. (°C) k f (°C/m) Acetic acid16.603.90 Benzene5.514.90 Naphthalene80.26.8 Water0.001.86

40. Benzophenone freezes at 48.1  C. A solution of 1.05 g urea ((NH 2 ) 2 CO, molar mass = 60.06 g/mol) in 30.0 g of benzophenone freezes at 42.4  C. What is k f for benzophenone? Example: Calculate k f

41. Benzophenone freezes at 48.1  C and has a k f of 9.8  C/molal. A 2.50-g sample of solute whose molar mass is 130.0 g/mol is dissolved in 32.0 g of benzophenone. What is the freezing point of the solution? Test Your Skill

42. Semipermeable membranes allow water and small molecules to pass through them. Osmosis is the diffusion of a fluid through a semipermeable membrane. Osmosis

43. When a semipermeable membrane separates a solution from the pure solvent, the net effect is for pure solvent to move through the membrane into the solution. The higher level of liquid produces an additional pressure, called osmotic pressure. Osmosis (continued)

44. Osmotic pressure is a colligative property, and can be calculated by the equation  = MRT where:  = osmotic pressure M = molar concentration of solute R = ideal gas law constant T = temperature in Kelvin Osmotic Pressure

45. A 5.70 mg sample of protein is dissolved in water to give 1.00 mL of solution. Calculate the molar mass of the protein if the solution has an osmotic pressure of 6.52 torr at 20  C. Example: Molar Mass by Osmotic Pressure

46. Colligative Properties - Summary PropertySymbolConc. UnitConstant Vapor pressure ∆P∆PMole fractionP°P° Boiling Point∆Tb∆Tb MolalKbKb Freezing point ∆Tf∆Tf MolalKfKf Osmotic pressure Π MolarRT

47. Electrolyte Solutions The colligative properties of electrolyte solutions are more pronounced because electrolytes separate into ions in solution. The van’t Hoff factor, i, is defined by the equation

48. The van’t Hoff Factor In dilute solution, the van’t Hoff factor for salts approaches the number of ions produced by one formula unit of the substance. NaCl → Na + (aq) + Cl - (aq)i = 2 MgBr 2 → Mg 2+ (aq) + 2Br - (aq)i = 3 The van’t Hoff factor generally decreases as the concentration increases.

49. Example: the van’t Hoff Factor Arrange the following aqueous solutions in order of increasing boiling points: 0.03 m urea (a nonelectrolyte), 0.01 m NaOH, 0.02 m BaCl 2, 0.01 m Fe(NO 3 ) 3.

50. Mixtures of Volatile Substances In a solution of two or more volatile compounds, all components of the mixture are in equilibrium with their vapors. An ideal solution is one in which all volatile components obey Raoult’s law for all compositions. P A =  A P  A, P B =  B P  B, P C =  C P  C, etc.

51. Ideal Solutions Mixtures of toluene and benzene form nearly ideal solutions.

52. At 27  C, the vapor pressure of carbon tetrachloride (CCl 4 ) is 127 torr and that of chloroform (CHCl 3 ) is 212 torr. What is the partial pressure of each substance, and the total vapor pressure of the solution, of a solution that contains 0.40 mol of CCl 4 and 0.60 mol of CHCl 3 ? Example: Vapor Pressure of Solutions

53. Distillation Distillation is the separation of a mixture of components based on differences in volatility (vapor pressure) by repeated evaporation and condensation of the mixture. The vapor always contains a larger mole fraction of the more volatile component.

54. Distillation Apparatus

55. Deviations from Raoult’s Law Most liquid-liquid solutions deviate from the ideal behavior predicted by Raoult’s law. Solutions have positive deviations if the vapor pressure is higher than predicted. Solutions have negative deviations if the vapor pressure is lower than predicted.