1. Chapter 11 Properties of solutions

2. Solutions A solution is a homogenous mixture. The solvent does the dissolving. The solute is dissolved by the solvent. There are various types solutions. Gas-gas: air, natural gas (gas solution) Liquid-liquid: antifreeze, vodka in water (liquid solution) Solid – solid: brass (solid solution) We will focus on aqueous solutions.

3. Solutions Gas-liquid: carbonated water (liquid solution) Solid – liquid: seawater, sugar water (liquid solution) Gas-solid: hydrogen in platinum (solid solution) We will focus on aqueous solutions.

4. Ways of Measuring Molarity = moles of solute Liters of solution Mass % = Mass of solute x 100 Mass of solution Mole fraction of component A

5. Molality = moles of solute Kilograms of solvent Molality is abbreviated m (lowercase) Normality - read but don’t focus on it. Note: molarity is based on volume so it changes slightly with temperature, molality does not because it depends on mass Ways of Measuring

6. Parts per million (ppm)- used for very dilute concentration, usually water or soil ppm = 1mg of substance 1 L H 2 O ppm = 1 mg of substance 1 kg of soil Ways of Measuring

7. Example A solution of phosphoric acid was made by dissolving 10.0 g of H 3 PO 4 in 100.0 mL of water. The resulting volume was 104.0 mL. Calculate the density, mole fraction, molarity, and molality of the solution. Assume water has a density of 1.00 g/mL.

8. Solution Formation

9. Heat of Solution The heat of solution is the amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent.

10. Heat of Solution Heat of solution is positive if energy is absorbed Heat of solution is negative if energy is released Equation for heat of solution

11. Steps in Solution Formation  H 1 Step 1 -  H 1 Step 1 - Expanding the solute Separating the solute into individual components

12. Steps in Solution Formation  H 2 Step 2 -  H 2 Step 2 - Expanding the solvent Overcoming intermolecular forces of the solvent molecules

13. Steps in Solution Formation  H 3 Step 3 -  H 3 Step 3 - Interaction of solute and solvent to form the solution

14. Predicting Solution Formation

15. Factors affecting Solubility

16. Structure Effects Molecular structure determines polarity Polarity similarities between solute and solvent favors solubility SO….there is a connection between structure and solubility

17. Structure Effects Consider vitamin A and vitamin C… Vitamin A is fat-soluble because it is nonpolar Vitamin C is polar

18. Copyright © Houghton Mifflin Company. All rights reserved.11–18 Figure 11.4 The Molecular Structures of (a) Vitamin A and (b) Vitamin C

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

20. Pressure effects Pressure affects the amount of gas that can dissolve in a liquid. Solubility of gases increase with pressure

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

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

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

24. Henry’s Law C=kP P = pressure of the gas above the liquid (in atm) C = concentration of the dissolved gas in (in mol/L) k = Henry’s law constant (specific to the gas) (mol/L  atm) NOTE: only when there is no chemical reaction between the solute and solvent

25. Example A bottle of soda at 25  C contains CO 2 gas at a pressure of 5.0 atm over the liquid. Assuming the partial pressure of CO 2 in the atmosphere is 4.0 x 10 -4, calculate the equilibrium concentrations before and after the bottle is opened. (k= 3.1 x 10 -2 mol/L  atm.)

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

27. Temperature effects Solids dissolve more rapidly at higher temperatures The amount of most solids increase with temperature BUT not always. Determined experimentally

28. Solubility Chart

29. Temperature effects Solubility of gases increases as temperature decrease

30. Vapor Pressures of solutions

31. Vapor Pressure of Solutions What does the term “vapor pressure” mean? The pressure of the vapor over a liquid at equilibrium. What?!! I don’t get it! Can you explain it more? When the rate of condensation is equal to the rate of evaporation, the pressure above the liquid is called “VAPOR PRESSURE”

32. Vapor Pressure of Solutions What happens to the vapor pressure of a solvent if we add a nonvolatile solute? The vapor pressure of the solvent decreases. OK, I can accept that….but can you tell me WHY? Because the dissolved nonvolatile solute decreases the number of solvent molecules AND so there are not as many solvent molecules that can escape

33. Vapor Pressure of Solutions Raoult observed this to be true….and related them together through P solution =  solvent P 0 solvent What happens when both the solute and solvent are volatile- (liquid-liquid solutions)? We modify Raoult’s law to P total = P solute + P solvent =  solute P 0 solute +  solvent P 0 solvent

34. Ideal solutions The application of Raoult’s law works for an ideal solution. So what is an IDEAL solution?!!! A solution where the solute and solvent interactions are similar.

35. Nonideal solutions What makes a solution nonideal? If hydrogen bonding occurs When ∆H soln is large and negative Both cause a negative deviation from Raoult’s law (lower vapor pressure) Can you explain this? Both of these cause interactions between the solute and solvent. So less molecules will escape and result in a lower vapor pressure.

36. Nonideal solutions What if ∆H soln is positive, what kind of deviation would occur? Cause a positive deviation from Raoult’s law (higher vapor pressure)

37. Colligative Properties, Osmotic Pressure and Electrolytes

38. Colligative Properties What are the colligative properties? Boiling point elevation Freezing point depression Osmotic pressure What solute characteristic do these have in common? Only dependent on the NUMBER of solute particles in an ideal solution

39. Boiling point Elevation If atmospheric pressure is 1 atm, what is the vapor pressure when a liquid boils? The vapor pressure of the liquid must be..1 atm! In order for a liquid to boil, vapor and atmospheric pressure must be equal. How would a nonvolatile solute affect vapor pressure? It lowers it. How does this affect the boiling point? It raises it. We need more energy (heat) to raise the vapor pressure.

40. Boiling point Elevation The equation is:  T = K b m solute  T is the change in the boiling point K b is a constant determined by the solvent. m solute is the molality of the solute I know you can use this equation to calculate the elevated boiling point…. We can also calculate the molar mass of the solute using this equation…YIPEE!

41. Boiling point Elevation A 2.00 gram sample of a biomolecule was dissolved in 15.0 g of CCl 4. The boiling point of this solution was determined to be 77.85 o C. Calculate the molar mass of this biomolecule? The boiling point of pure CCl 4 is 76.50 o C and K b for CCl 4 is 5.03 o C ·kg/mol

42. Freezing point Depression Why will a nonvolatile solute dissolved in a liquid lower its freezing point? Use water as the liquid in your explanation. The solute causes the vapor pressure of water to be lower than ice. In order for the vapor pressure of ice = the vapor pressure of liquid, the temperature must decrease.

43. Freezing point Depression The equation is:  T = -K f m solute  T is the change in the freezing point K f is a constant determined by the solvent m solute is the molality of the solute Hey! We can also use this to calculate molar mass of the solute! COOL!

44. Freezing point Depression When 0.455 g of thyroxine is dissolved in 10.0 g benzene, the freezing point of the solution is depressed by 0.300 o C. What is the molar mass of thyroxine? K f for benzene = 5.12 O C · kg/mol

45. Osmotic Pressure What is osmosis? Flow of solvent through a semipermeable membrane Semipermeable membrane! What so special about that? It’s special because it allows solvent molecules to pass through but blocks solute molecules

46. Osmotic Pressure What does this have to do with pressure? When the system has reached equilibrium, liquid levels are different. So the pressure on the solution is greater than that of the pure solvent. We call this OSMOTIC PRESSURE

47. Osmotic Pressure π=MRT ∏ is the osmotic pressure M is the molarity of the solution T is the temperature in Kelvin What do you think we can do with osmotic pressure? Yep! We can calculate the molar mass of a solute! Wow! Colligative proerties are useful!

48. Electrolytes in solution Colligative properties depend on the number of particles. So….. Would electrolytes have a different affect on colligative properties than non electrolytes? You betcha they would! When electrolytes dissolve they dissociate. This creates more particles. 1 mole of NaCl makes 2 moles of ions. 1mole Al(NO 3 ) 3 makes 4 moles ions.

49. Electrolytes in solution How many moles of particles would sucrose make? Why it would only have 1 mole of particles since its nonelectrolyte!

50. Electrolytes have a bigger impact 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 The expected value can be determined from the formula of the compound. We can change our formulas to  = iKm and ∏= i MRT Electrolytes in solution

51. The actual value is usually less because At any given instant some of the ions in solution will be paired up. Ion pairing increases with concentration. i decreases with increasing concentration. Expected vs Observed van’t Hoff Factors