Solutions

1. Solutions A homogeneous mixture of two or more substances.

2. The Solution Process We will focus on solid or liquid solutes dissolved in a liquid solvent. Since all particles are in contact with each other, the solute-solute and solvent-solvent forces of attraction are disrupted, and new, solute-solvent forces of attraction are created.

3. The Solution Process The disruption of solute-solute and solvent-solvent forces of attraction requires energy, and is endothermic. The interaction of solvent and solute usually releases energy. The sum of the energy of all three steps is called the enthalpy of solution, ΔHosoln. Note that solutions may form whether the net process is endothermic or exothermic.

4. The Solution Process

5. The Solution Process In addition to the enthalpy of solution, we must also consider the entropy of mixing. Entropy is a measure of randomness or disorder. An increase in entropy makes a process more likely to occur. Since mixing pure substances increases entropy, this factor makes processes that are slightly endothermic favorable.

6. Entropy of Mixing

7. The Solution Process The general rule on solution formation is: Like dissolves like. Polar and ionic compounds dissolve in polar solvents. Non-polar compounds dissolve in non-polar solvents.

8. Like Dissolves Like Vitamin A consists almost entirely of carbon and hydrogen, and is non-polar. As a result, vitamin A is fat-soluble, and can be stored in the body.

9. Like Dissolves Like Vitamin C contains polar C-O and O-H bonds. It is water soluble, and must be consumed often, as it is excreted easily. C-O bond O-H bonds

10. Like Dissolves Like Ionic Compounds

11. Like Dissolves Like

12. The Solution Process Disrupt-ion of solute Disrupt-ion of solvent Solute/Solvent interact-ion

13. Ionic Aqueous Solutions When an ionic compound is dissolved in water, the energy required to separate the ions of the solute is equal to –(lattice energy), or -ΔHlattice. The energy released as the gaseous ions dissolve in water is called the hydration energy, ΔHhydration. The net energy change is ΔHsoln.

14. Heat of Hydration

15. Factors Affecting Solubility • Molecular Structure • Pressure (for gaseous solutes) • Temperature

16. Pressure Effects Gases dissolved in a liquid solute obey Henry’s Law: C = kP where C is the concentration, k is a constant specific to solute and solvent, and P is the pressure of the gas above the solution

17. Pressure Effects Gases dissolved in a liquid solute obey Henry’s Law: C = kP The amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.

18. Henry’s Law

19. Pressure Effects

20. Temperature Effects For gases dissolved in liquids, the solubility decreases as temperature increases. That is, gases dissolve better in cold liquids than in warmer liquids.

21. Temperature Effects For solid solutes dissolved in water, the effect of temperature on solubility is difficult to predict, although many solids dissolve more as temperature increases.

22. Solution Concentration Although molarity (M) is used for stoichiometry calculations, there are many other ways to express the concentration of a solution. Molarity will vary slightly with changes in temperature as the volume of the solution expands or contracts. Units such as mass percent, mole fraction, or molality remain constant as temperature changes.

23. Solution Concentration Mass percent = (mass of solute) (100%) (mass of solution) Mole fraction (XA) = (moles of A) total # of moles Molality (m) = moles of solute kg of solvent

24. Very Dilute Solutions The concentration of very dilute solutions are expressed in parts per million (ppm) or parts per billion (ppb). ppm = [(mass solute) x 106 ] ÷(mass soln) ppb = [(mass solute) x 109 ] ÷(mass soln)

25. The Colligative Properties The colligative properties are properties that depend upon the concentration of particles (molecules or ions) dissolved in a volatile solvent, and not on the nature of the particles. They include: • vapor pressure • freezing point • boiling point • osmotic pressure

26. The Colligative Properties Relatively simple mathematical relationships can be used to predict the changes in vapor pressure, freezing and boiling point, etc. The properties can be predicted for dilute solutions (<0.1M) of non-volatile solute (usually solids) dissolved in a volatile solvent (usually a liquid).

27. Vapor Pressure The addition of a non-volatile solute to a volatile solvent lowers the vapor pressure of the solvent.

29. Vapor Pressure The decrease in vapor pressure can be understood by looking at the evaporation process. We need to compare the enthalpy change (ΔHvap) and entropy change of evaporation.

30. Vapor Pressure The vapor pressure of the pure solvent or the solution is the result of solvent molecules escaping the liquid surface and becoming gaseous. Since the solute is non-volatile, it does not evaporate. Since only solvent molecules evaporate, the enthalpy change for pure solvent or the solution is the same.

31. Vapor Pressure The decrease in vapor pressure of the solution is the result of changes in entropy. The vapor in either container is disordered, due to the random motion of gaseous solvent.

32. Vapor Pressure The liquid phases differ in entropy. The pure solvent is relatively ordered since all of the molecules are the same (solvent).

33. Vapor Pressure The liquid phase of the solution is much more random, since it is a mixture.

34. Vapor Pressure Upon evaporation, the pure solvent undergoes a greater increase in entropy than the solution.

35. Vapor Pressure Systems tend to maximize entropy. The pure solvent evaporates more readily, because it undergoes a greater increase in entropy.

36. Boiling Point Elevation

37. Vapor Pressure Lowering The change in vapor pressure can be calculated as follows: ∆vp = -Xsolute Psolvent where X is the mole fraction of solute particles Posolvent is the vapor pressure of the pure solvent o

38. Vapor Pressure Lowering ∆vp = -Xsolute Posolvent The sign is negative because the vapor pressure decreases.

39. Vapor Pressure Lowering Psoln = Xsolvent Posolvent The mole fraction of solvent, Xsolvent , = moles of solvent/total moles of particles and solvent.

40. Problem – Vapor Pressure Water has a vapor pressure of 92.6 mmHg at 50oC. a) Compare the vapor pressure of two aqueous solutions at 50oC. One contains .100 mole of sucrose dissolved in 1.00 mol of water. The other contains .100 moles of CaCl2 dissolved in 1.00 mol of water. b) Calculate the vapor pressure of the CaCl2 solution.

41. Solution Phase Diagrams The lowering of the vapor pressure due to the presence of a non-volatile solute affects several properties. The phase diagram for the solution will be shifted, due to the lower vapor pressure of the solution.

42. Solution Phase Diagrams

43. Solution Phase Diagrams As a result of the lower vapor pressure, the boiling point of the solution is greater than that of pure solvent.

44. Solution Phase Diagrams Since the liquid-solid line is shifted to a lower temperature, the freezing point of the solution is lowered.

45. Properties of Solutions Solutions of non-volatile solutes in a volatile solvent have - higher boiling points and • lower freezing points than the pure solvent.

46. Boiling Point Elevation The size of the increase in boiling point depends upon the concentration of solute particles. ∆Tb = Kbm(i) where Kb is the solvent dependent boiling point elevation constant, m = molality of the solute i = van’t Hoff factor

47. The van’t Hoff Factor, i The van’t Hoff factor is the number of particles in solution compared to the number dissolved. If an ionic compound forms two ions per formula unit, its i value = 2.

48. The van’t Hoff Factor, i If a molecule “pairs up” in solution, with two molecules uniting to form one molecule, then the i factor will be 0.5. For non-electrolytes, the i factor is usually 1, and is often ignored.

49. Freezing Point Depression The size of the decrease in freezing point depends upon the concentration of solute particles. ∆Tf = -Kfm(i) where Kf is the solvent dependent freezing point depression constant, m = molality of the solute i = van’t Hoff factor

50. Constants for Common Solvents