Chapter 13 Properties of Solutions

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

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

3. Solutions The ___________ forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

4. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and ___________, or solvates, them.

5. How Does a Solution Form If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the ___________ energy of the salt crystal.

6. Energy Changes in Solution • Simply put, three processes affect the ___________ of the process: • ___________ of solute particles • ___________ of solvent particles • New interactions between solute and solvent

7. Energy Changes in Solution The ___________ change of the overall process depends on H for each of these steps. Together, this is called the energy of solution formation or energy of “solvation”

8. Why Do Endothermic Processes Occur? Things do not tend to occur ___________ (i.e., without outside intervention) unless the energy of the system is lowered.

9. Why Do Endothermic Processes Occur? Yet we know that in some processes, like the dissolution of NH4NO3 in water, heat is ___________, not released.

10. Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as ___________) of a system tends to lower the energy of the system.

11. Enthalpy Is Only Part of the Picture So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more ___________.

12. Student, Beware! Just because a substance disappears when it comes in contact with a solvent, it ___________ mean the substance dissolved.

13. Student, Beware! • 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 ___________.

14. Types of Solutions • Saturated • Solvent holds as much ___________ as is possible at that temperature. • Dissolved solute is in dynamic ___________ with solid solute particles.

15. Types of Solutions • Unsaturated • Less than the ___________ amount of solute for that temperature is dissolved in the solvent.

16. Types of Solutions • Supersaturated • Solvent holds ___________ solute than is normally possible at that temperature. • These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

17. Factors Affecting Solubility • Chemists use the axiom “like ___________ like”: • ___________ substances tend to dissolve in polar solvents. • ___________ substances tend to dissolve in nonpolar solvents.

18. Factors Affecting Solubility The more similar the ___________ attractions, the more likely one substance is to be soluble in another.

19. Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has ___________ forces) is not.

20. Factors Affecting Solubility • Vitamin A is soluble in ___________ compounds (like fats). • Vitamin C is soluble in ___________. Other vitamins behave one way or the other…limey story…

21. Gases in Solution • In general, the solubility of gases in water increases with ___________ mass. • Larger molecules have stronger ___________ forces.

22. Gases in Solution • The solubility of liquids and solids does not change appreciably with ___________. • The solubility of a gas in a liquid is ___________ proportional to its pressure.

23. 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 ___________ of the gas above the liquid.

24. Temperature Generally, the ___________ of solid solutes in liquid solvents increases with increasing temperature.

25. Temperature • The opposite is true of gases: • Carbonated soft drinks are more “___________” if stored in the refrigerator. • Warm lakes have less O2___________ in them than cool lakes.

26. Ways of Expressing Concentrations of Solutions

27. mass of A in solution ___________ mass of solution Mass Percentage Mass % of A =  100

28. 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 =

29. moles of A __________ 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!

30. 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 ___________.

31. mol of solute kg of solvent m = Molality (m) Because both moles and mass do not change with temperature, molality (unlike molarity) is ___________ temperature dependent.

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

33. Colligative Properties • Changes in ___________ properties depend only on the ___________ of solute particles present, not on the ___________ of the solute particles. • Among colligative properties are • Vapor pressure lowering • Boiling point elevation • Melting point depression • Osmotic pressure

34. Vapor Pressure Because of solute-solvent intermolecular attraction, higher concentrations of ___________ solutes make it harder for solvent to escape to the vapor phase.

35. Vapor Pressure Therefore, the vapor pressure of a solution is ___________ than that of the pure solvent. Check out the boiling point of salt water vs. “regular water.”

36. 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.

37. Boiling Point Elevation and Freezing Point Depression Nonvolatile solute-solvent interactions also cause solutions to have ___________ boiling points and ___________ freezing points than the pure solvent.

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

39. Example problems

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

41. Note that 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

42. Colligative Properties of Electrolytes Since these properties depend on the number of particles dissolved, solutions of ___________ (which dissociate in solution) should show greater changes than those of nonelectrolytes. (NaCl --> vs. sugar -->)

43. 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.

44. Write the “dissolving equations” and count the 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 for a short time, 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 ___________ dependent.

48. The van’t Hoff Factor We modify the previous equations by ___________ by the van’t Hoff factor, i Tf = Kf •m•i

49. Osmosis • Some substances form ______________ membranes, allowing some smaller particles to pass through, but blocking other larger particles. • In biological systems, most semipermeable membranes allow water to pass through, but ______________ are not free to do so.

50. Osmosis In osmosis, there is net movement of solvent from the area of ______________ solvent concentration (lowersoluteconcentration) to the are of ______________ solvent concentration (highersoluteconcentration).