Chapter 13 Properties of Solutions

2. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 13Properties of Solutions Adapted by SA Green from: John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

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

4. Solutions How does a solid dissolve into a liquid? What ‘drives’ the dissolution process? What are the energetics of dissolution?

5. How Does a Solution Form? • Solvent molecules attracted to surface ions. • Each ion is surrounded by solvent molecules. • Enthalpy (DH) changes with each interaction broken or formed. Ionic solid dissolving in water

6. How Does a Solution Form? • Solvent molecules attracted to surface ions. • Each ion is surrounded by solvent molecules. • Enthalpy (DH) changes with each interaction broken or formed.

7. How Does a Solution Form The ions are solvated (surrounded by solvent). If the solvent is water, the ions are hydrated. The intermolecular force here is ion-dipole.

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

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

10. Enthalpy changes during dissolution DHsoln = DH1 + DH2 + DH3 The enthalpy of solution, DHsoln, can be either positive or negative. DHsoln (MgSO4)= -91.2 kJ/mol --> exothermic DHsoln (NH4NO3)= 26.4 kJ/mol --> endothermic

11. Why do endothermic processes sometimes occur spontaneously? Some processes, like the dissolution of NH4NO3 in water, are spontaneous at room temperature even though heat is absorbed, not released.

12. Enthalpy Is Only Part of the Picture Entropy is a measure of: • Dispersal of energy in the system. • Number of microstates (arrangements) in the system. b. has greater entropy,  is the favored state (more on this in chap 19)

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

14. SAMPLE EXERCISE 13.1Assessing Entropy Change In the process illustrated below, water vapor reacts with excess solid sodium sulfate to form the hydrated form of the salt. The chemical reaction is Does the entropy of the system increase or decrease?

15. dry Dissolution vs reaction NiCl2(s) Ni(s) + HCl(aq) NiCl2(aq) + H2(g) • Dissolution is a physical change—you can get back the original solute by evaporating the solvent. • If you can’t, the substance didn’t dissolve, it reacted.

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 • Unsaturated Solution • Less than the maximum amount of solute for that temperature is dissolved in the solvent. • No solid remains in flask.

18. 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.

19. Degree of saturation Unsaturated, Saturated or Supersaturated?  How much solute can be dissolved in a solution? More on this in Chap 17 (solubility products, p 739)

20. Factors Affecting Solubility • Chemists use the axiom “like dissolves like”: • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents.

21. Factors Affecting Solubility The stronger the intermolecular attractions between solute and solvent, the more likely the solute will dissolve. Example: ethanol in water Ethanol = CH3CH2OH Intermolecular forces = H-bonds; dipole-dipole; dispersion Ions in water also have ion-dipole forces.

22. Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water. Cyclohexane (which only has dispersion forces) is not water-soluble.

23. Factors Affecting Solubility • Vitamin A is soluble in nonpolar compounds (like fats). • Vitamin C is soluble in water.

24. Which vitamin is water-soluble and which is fat-soluble?

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

26. Gases in Solution

27. 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.

28. Henry’s Law Sg = kPg where • Sg is the solubility of the gas; • k is the Henry’s law constant for that gas in that solvent; • Pg is the partial pressure of the gas above the liquid.

29. Henry’s Law k for N2 at 25° =6.8 x 10-4 mol/L atm Sg = kPg

30. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

31. 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 O2 dissolved in them than cool lakes.

32. Chap 13:Ways of Expressing Concentrations of Solutions

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

34. mass of A in solution total mass of solution mass of A in solution total mass of solution Parts per Million andParts per Billion Parts per Million (ppm) ppm =  106 Parts per Billion (ppb)  109 ppb =

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

36. mol of solute L of solution M = Molarity (M) • You will recall this concentration measure from Chapter 4. • Because volume is temperature dependent, molarity can change with temperature.

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

38. Mass/Mass Moles/Mass Moles/Moles Moles/L

39. SAMPLE EXERCISE 13.4 Calculation of Mass-Related Concentrations (a) A solution is made by dissolving 13.5 g of glucose (C6H12O6) in 0.100 kg of water. What is the mass percentage of solute in this solution? (b) A 2.5-g sample of groundwater was found to contain 5.4g of Zn2+ What is the concentration of Zn2+ in parts per million? PRACTICE EXERCISE (a) Calculate the mass percentage of NaCl in a solution containing 1.50 g of NaCl in 50.0 g of water. (b) A commercial bleaching solution contains 3.62 mass % sodium hypochlorite, NaOCl. What is the mass of NaOCl in a bottle containing 2500 g of bleaching solution? PRACTICE EXERCISE A commercial bleach solution contains 3.62 mass % NaOCl in water. Calculate (a) the molality and (b) the mole fraction of NaOCl in the solution.

40. 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

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

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

43. Raoult’s Law PA = XAPA where • XA is the mole fraction of compound A • PA is the normal vapor pressure of A at that temperature NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.

44. SAMPLE EXERCISE 13.8 Calculation of Vapor-Pressure Lowering Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25°C. Calculate the vapor pressure at 25°C of a solution made by adding 50.0 mL of glycerin to 500.0 mL of water. The vapor pressure of pure water at 25°C is 23.8 torr (Appendix B). PRACTICE EXERCISE The vapor pressure of pure water at 110°C is 1070 torr. A solution of ethylene glycol and water has a vapor pressure of 1.00 atm at 110°C. Assuming that Raoult’s law is obeyed, what is the mole fraction of ethylene glycol in the solution?

45. 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.

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

47. Freezing Point Depression • The change in freezing point can be found similarly: Tf = Kf  m • Here Kf is the molal freezing point depression constant of the solvent. Tf is subtracted from the normal freezing point of the solvent.

48. In both equations, T does not depend on what the solute is, but only on how many particles are dissolved. Tb = Kb  m Tf = Kf  m Boiling Point Elevation and Freezing Point Depression

49. 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.

50. 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.