Chapter 10:Solutions

1. Chapter 10:Solutions

2. Types of Solutions • Solid/Solid: • Steal • Brass • Metal solutions are called alloys • Liquid/Liquid: • Rubbing alcohol • Gas/Gas: • Air • Gas/Liquid • Carbonated sodas • Solid/Liquid • Salt in water • Liquid/Gas • Humidity (water in air)

3. Solutions • The solvent does the dissolving. • The solute is dissolved. • There are many types of solvents dissolving all types of solutes.

4. Ways of Measuring • Molarity (M) = moles of solute Liters of solution • % mass = Mass of solute x 100 Mass of solution • Mole fraction (X) of component A XA = moles solute A moles total solution • Molality (m)= moles of solute Kilograms of solvent

5. Solvation • Dissociation • separation of an ionic solid into aqueous ions NaCl(s)  Na+(aq) + Cl–(aq)

6. Solvation • Ionization • breaking apart of some polar molecules into aqueous ions HNO3(aq) + H2O(l)  H3O+(aq) + NO3–(aq)

7. B. Solvation • Molecular Solvation • molecules stay intact C6H12O6(s)  C6H12O6(aq)

8. SATURATED SOLUTION no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form UNSATURATED SOLUTION more solute dissolves C. Solubility concentration

9. C. Solubility • Solubility • maximum grams of solute that will dissolve in 100 g of solvent at a given temperature • varies with temp • based on a saturated soln

10. C. Solubility • Solubility Curve • shows the dependence of solubility on temperature

11. Dissolution of a Solute • Depends on 2 factors: • Change in energy (enthalpy) • Change in disorder (entropy) • Process is favored by: • Decrease in enthalpy (H) - exothermic • Increase in entropy (S)

12. Formation of a Solution • Heat of solution ( DHsoln ) is the energy change for making a solution. • - DHsoln = exothermic (favored) • Most easily understood if broken into steps. 1.Break apart solvent 2.Break apart solute 3. Mixing solvent and solute

13. Energy of Making Solutions 1. Break apart Solvent • Have to overcome attractive forces. DH1 >0 2. Break apart Solute. • Have to overcome attractive forces. DH2 >0

14. Energy of Making Solutions 3. Mixing solvent and solute • DH3 depends on what you are mixing. • Molecules that attract each other have a DH3 thatis large and negative. • Molecules that do not attract have a DH3 thatissmall and negative. • This explains the rule “Like dissolves Like”

15. Solute DH3 DH2 Solvent DH1 DH3 Solution B Solution A* Size of DH3 determines whether solution will form Energy Solute & Solvent * Solution A is less likely to form since the process requires increased energy, but may still occur due to increased entropy.

16. Formation of Solutions • If DHsoln is small and positive, a solution will still form because of increased entropy. • There are many more ways for them to become mixed than there is for them to stay separate.

17. Structure and Solubility • Water soluble molecules must have dipole moments -polar bonds. • To be soluble in non polar solvents the molecules must be non polar.

18. O- P O- CH2 CH2 CH2 CH3 CH2 O- CH2 CH2 CH2 Soap • Hydrophobic non-polar end • Hydrophilic polar end

19. Dissolution of Solutes in Liquids • Dissolution of Solids in Liquids • Solvation: the process by which solvent molecules surround and interact with solute ions or molecules • When solvation occurs with water as the solvent it is called hydration • DHsoln = (heat of solvation) – (crystall lattice energy)

20. Dissolution of Solutes in Liquids • Dissolution of Liquids in Liquids • Miscibility: the ability of one liquid to dissolve in another • Polar liquids tend to interact with polar solvents • Non polar liquids tend to interact with non polar solvents

21. Dissolution of Solutes in Liquids • Dissolution of Gases in Liquids • Polar gases tend to interact with polar solvents • Non polar gases tend to interact with non polar solvents

22. Rates of Dissolution & Saturation • Increasing surface area increases the rate of dissolution. dissolution Soliddissolved particles crystallization • Saturated = equilibrium is reached between the solid and dissolved ions in a solution • Supersaturated = when a solution contains more ions in the dissolved state than would normally occur; usually by dissolving the excess solute with heat then cooling slowly

23. Effect of Temperature on Solubility • For solutes that go into solution endothermically (require energy) solubility increases when the temperature increases. H2O KCl + 17.2 kJ --- K+ (aq) + Cl-(aq) • Increasing the temperature would increase the solubility for KCl

24. Effect of Temperature on Solubility • For solutes that go into solution exothermically (release energy) solubility decreases when the temperature increases. • Oxygen’s solubility decreases from 0.041 grams/liter at 25oC to only 0.027 grams/liter at 50oC.

25. Temperature & Solubility of Gases • As temperature increases, solubility decreases. • Gas molecules can move fast enough to escape.

26. Pressure effects • Changing the pressure doesn’t effect the amount of solid or liquid that dissolves • Solids and liquids are incompressible. • Pressure does effect the solubility gases.

27. The dissolved gas and the gas above the liquid are at equilibrium • The equilibrium is dynamic.

28. If you increase the pressure the gas molecules dissolve faster. • The equilibrium is disturbed.

29. The system reaches a new equilibrium with more gas dissolved.

30. Henry’s Law • Henry’s Law. Pgas= kCgas Pgas = pressure of gas above the solution *k = a constant for a particular gas at a particular temperature Cgas = concentration of the gas in solution expressed as molarity or mole fraction

31. Colligative Properties • Colligative Properties: physical properties of solutions that depend on the number (not kind) of solute particles in a given amount of solvent. • Vapor pressure lowering • Boiling point elevation • Freezing point depression • Osmotic pressure

32. Vapor Pressure of Solutions • A nonvolatile liquid lowers the vapor pressure of the solution. • A solid solute also lowers the vapor pressure. • The molecules of the solventmust overcome the force of both the other solvent molecules and the solute molecules.

33. Why does Vapor Pressure change? • When a solute is dissolved in a liquid, some of the volume is occupied by the solute particles. • There are fewer solvent molecules per unit of area at the surface. • In a pure solvent molecules vaporize to increase entropy (disorder). • In a solution there is already an increase in disorder so vaporization is slowed. Note: Solutions of gases and volatile liquid solutes have higher vapor pressures than their respective pure solvents so this colligative property does not apply.

34. Raoult’s Law: • Psoln = Xsolvent x Posolvent Psoln = Vapor pressure of the solution Xsolvent =mole fraction of the solvent Posolvent = vapor pressure of the pure solvent • Applies only to an ideal solution where the solute doesn’t contribute to the vapor pressure.

35. Raoult’s Law Practice Problem

36. Calculating Molar Mass from A Change in Vapor Pressure • The equation: Psoln = Xsolvent x Posolvent • Can be rewritten in a slightly different form to suggest another useful technique. The mole fraction of solvent is given as; Xsolvent = moles solvent moles solvent + moles solute

37. Calculating Molar Mass from A Change in Vapor Pressure • Which can be further transformed to; Xsolvent = (grams/MM)solvent (grams/MM)solvent + (grams/MM)solute • From this relationship it is possible to determine the molar mass of a compound from the change in vapor pressure of the solvent.

38. Calculating Molar Mass from A Change in Vapor Pressure

39. Colligative Properties • Because dissolved particles affect vapor pressure - they affect phase changes. • Useful for determining molar mass

40. Boiling point Elevation • Because a non-volatile solute lowers the vapor pressure it raises the boiling point. • The equation is: DT = Kbmsolute • DT is the change in the boiling point • Kb is a constant determined by the solvent. • msolute is the molality of the solute

41. Freezing point Depression • Because a non-volatile solute lowers the vapor pressure of the solution it lowers the freezing point. • The equation is: DT = Kfmsolute • DT is the change in the freezing point • Kf is a constant determined by the solvent • msolute is the molality of the solute

42. Calculations t: change in temperature (°C) k: constant based on the solvent (°C·kg/mol) m: molality (m) d.f.: dissocation factor (# of particles) t = m x d.f. x k

43. Calculations • Dissociation Factor (# of Particles) • Nonelectrolytes (covalent) • remain intact when dissolved • 1 particle • Electrolytes (ionic) • dissociate into ions when dissolved • 2 or more particles

44. Sample Freezing Point & Boiling Point Problem

45. Calculating Molar Mass from Boiling Point Elevation or Freezing Point Depression • We can also use the expression to determine the molecular weight of an unknown solute. If we recognize that; • and that;

46. Calculating Molar Mass from Boiling Point Elevation or Freezing Point Depression • substituting into the freezing point expression we have; • rearranging and solving for MW we have;

47. Calculating Molar Mass from Boiling Point Elevation or Freezing Point Depression

48. Electrolytes in solution • Since colligative properties only depend on the number of molecules. • Ionic compounds should have a bigger effect. • When they dissolve they dissociate. • Individual Na and Cl ions fall apart. • 1 mole of NaCl makes 2 moles of ions. • 1mole Al(NO3)3 makes 4 moles ions.

49. Electrolytes in solution • Electrolytes have a bigger impact on on melting and freezing points per mole because they make more pieces. • Relationship is expressed using the van’t Hoff factor(i) i = meffective mstated • The expected value can be determined from the formula.

50. Electrolytes in solution • The actual value is usually less because • At any given instant some of the ions in solution will be paired. • Ion pairing increases with concentration. • i decreases with in concentration.