1. Solutions

2. Occur in all phases u The solvent does the dissolving. u The solute is dissolved. u There are examples of all types of solvents dissolving all types of solvent. u We will focus on aqueous solutions.

3. Ways of Measuring u Molarity = moles of solute Liters of solvent u % mass = Mass of solute x 100 Mass of solution  Mole fraction of component A  A = N A N A + N B

4. u Molality = moles of solute Kilograms of solvent  Molality is abbreviated m u Normality - read but don’t focus on it. Ways of Measuring

5. Energy of Making Solutions  Heat of solution (  H soln ) is the energy change for making a solution. u Most easily understood if broken into steps. u 1.Break apart solvent u 2.Break apart solute u 3. Mixing solvent and solute

6. 1. Break apart Solvent  Have to overcome attractive forces.  H 1 >0 2. Break apart Solute.  Have to overcome attractive forces.  H 2 >0

7. 3. Mixing solvent and solute   H 3 depends on what you are mixing.  Molecules can attract each other  H 3 is large and negative.  Molecules can’t attract-  H 3 is small and negative. u This explains the rule “Like dissolves Like”

8. EnergyEnergy Reactants Solution H1H1 H2H2 H3H3 Solvent Solute and Solvent  Size of  H 3 determines whether a solution will form H3H3 Solution

9. Types of Solvent and solutes  If  H soln is small and positive, a solution will still form because of entropy. u There are many more ways for them to become mixed than there is for them to stay separate.

10. Structure and Solubility u Water soluble molecules must have dipole moments -polar bonds. u To be soluble in non polar solvents the molecules must be non polar. u Read Vitamin A Vitamin C discussion pg. 509

11. Soap PO-O- CH 3 CH 2 O-O- O-O-

12. Soap u Hydrophobic non- polar end PO-O- CH 3 CH 2 O-O- O-O-

13. Soap u Hydrophilic polar end PO-O- CH 3 CH 2 O-O- O-O-

14. PO-O- CH 3 CH 2 O-O- O-O- _

15. u A drop of grease in water u Grease is non-polar u Water is polar u Soap lets you dissolve the non-polar in the polar.

16. Hydrophobic ends dissolve in grease

17. Hydrophilic ends dissolve in water

18. u Water molecules can surround and dissolve grease. u Helps get grease out of your way.

19. Pressure effects u Changing the pressure doesn’t effect the amount of solid or liquid that dissolves u They are incompressible. u It does effect gases.

20. Dissolving Gases u Pressure effects the amount of gas that can dissolve in a liquid. u The dissolved gas is at equilibrium with the gas above the liquid.

21. u The gas is at equilibrium with the dissolved gas in this solution. u The equilibrium is dynamic.

22. u If you increase the pressure the gas molecules dissolve faster. u The equilibrium is disturbed.

23. u The system reaches a new equilibrium with more gas dissolved. u Henry’s Law. P= kC Pressure = constant x Concentration of gas

24. Temperature Effects u Increased temperature increases the rate at which a solid dissolves. u We can’t predict whether it will increase the amount of solid that dissolves. u We must read it from a graph of experimental data.

25. 20 406080 10 0

26. Gases are predictable u As temperature increases, solubility decreases. u Gas molecules can move fast enough to escape. u Thermal pollution.

27. Vapor Pressure of Solutions u A nonvolatile solvent lowers the vapor pressure of the solution. u The molecules of the solvent must overcome the force of both the other solvent molecules and the solute molecules.

28. Raoult’s Law:  P soln =  solvent x P solvent u Vapor pressure of the solution = mole fraction of solvent x vapor pressure of the pure solvent u Applies only to an ideal solution where the solute doesn’t contribute to the vapor pressure.

29. Aqueous Solution Pure water u Water has a higher vapor pressure than a solution

30. Aqueous Solution Pure water u Water evaporates faster from for water than solution

31. u The water condenses faster in the solution so it should all end up there. Aqueous Solution Pure water

32. Colligative Properties u Because dissolved particles affect vapor pressure - they affect phase changes. u Colligative properties depend only on the number - not the kind of solute particles present u Useful for determining molar mass

33. Boiling point Elevation u Because a non-volatile solute lowers the vapor pressure it raises the boiling point.  The equation is:  T = K b m solute   T is the change in the boiling point u K b is a constant determined by the solvent.  m solute is the molality of the solute

34. Freezing point Depression u Because a non-volatile solute lowers the vapor pressure of the solution it lowers the freezing point.  The equation is:  T = K f m solute   T is the change in the freezing point u K f is a constant determined by the solvent  m solute is the molality of the solute

35. 1 atm Vapor Pressure of solution Vapor Pressure of pure water

36. 1 atm Freezing and boiling points of water

37. 1 atm Freezing and boiling points of solution

38. 1 atm TfTf TbTb

39. Electrolytes in solution u Since colligative properties only depend on the number of molecules. u Ionic compounds should have a bigger effect. u When they dissolve they dissociate. u Individual Na and Cl ions fall apart. u 1 mole of NaCl makes 2 moles of ions. u 1mole Al(NO 3 ) 3 makes 4 moles ions.

40. u 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 = Moles of particles in solution Moles of solute dissolved u The expected value can be determined from the formula.

41. u The actual value is usually less because u At any given instant some of the ions in solution will be paired. u Ion pairing increases with concentration.  i decreases with in concentration. u We can change our formulas to  H = iKm

42. u Label your solutions, in the flasks and the ice cube trays. u Final conclusion will be to compare the actual freezing point depression to the theoretical. u Give reasons for any differences.

43. u Liquid-liquid solutions where both are volatile. u Modify Raoult’s Law to  P total = P A + P B =  A P A 0 +  A P B 0 u P total = vapor pressure of mixture u P A 0 = vapor pressure of pure A u If this equation works then the solution is ideal. u Solvent and solute are alike. Ideal solutions

44. Deviations u If Solvent has a strong affinity for solute (H bonding). u Lowers solvents ability to escape. u Lower vapor pressure than expected. u Negative deviation from Raoult’s law.   H soln is large and negative exothermic.  Endothermic  H soln indicates positive deviation.