Solutions

1. Solutions Chapters 4.1, 4.5, and 13

2. Solutions Solutions are homogeneous mixtures. Solutions are substances composed of dissolved particles(solute) in a dissolving medium(solvent).

3. Electrolytic Properties • Electrolytes – when ions are dissolved into an aqueous solution they can conduct an electric current; the more ions it breaks down into, the stronger electrolyte it is. • Nonelectrolytes – a substance that doesn’t form ions in solution, thus it doesn’t conduct an electric current.

4. Strong Electrolytes = completely ionized when dissolved in water. • Strong Acids and Strong Bases are Strong Electrolytes. • Weak Electrolytes = slightly ionized solution when dissolved in water • Weak acids and bases are weak electrolytes.

5. Ionic vs Molecular Compounds in Water • Ionic compoundsbreak down into their component ions in water. Na2SO4 becomes Na+ ions and SO4-2 ions dispursed through the water. • Molecular compounds dissolve in water, but the entire molecule stays intact. C6H12O6 dissolves and is dispersed through water, but the entire C6H12O6 molecule remains intact.

6. Concentrations of Solutions • Molarity = Moles of solute / Liters of Sol’n • Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125 mL of solution.

7. Expressing Conc. Of an Electrolyte • When ionic compounds dissolve, the relative concentrations depend on the chemical formula. • 1.0 M of NaCl is 1.0 M of Na+ and 1.0 M of Cl- • If I had 1.0 M of Na2SO4, what is the concentration of Na+ and SO4-2?

8. Preparation of Standard Solutions • Standard solutions = solution whose concentration is accurately known. • How much x How strong x What does it weigh? • L x mol / L x g / mol = grams required to prepare the standard

9. Dilutions • You can make substances less concentrated by adding water to the solution. • Moles of solute before dilution = moles of solute after dilution • M1V1 = M2V2 • How many mL of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?

10. NaCl dissolving in water.

11. The Solution Process • As individual solute ions break away from the crystal, the negatively and positively charged ions become surrounded by solvent molecules and the ionic crystal dissolves. • This is caused by Random Molecular Motion.

12. “Like dissolves Like” • Polar molecules and ionic compounds tend to dissolve in polar solvents. • Nonpolar molecules dissolve in nonpolar solvents.

13. Energy Changes and Solution Formation • Heat of solution ( DHsoln ) is the energy change for making a solution. • Most easily understood if broken into steps: • Break apart solute (endothermic) • Break apart solvent (endothermic) • Mixing solvent and solute (often exothermic)

14. 1. Break apart the Solute • Have to overcome IMFs to break the solute into individual components (expanding the solute). • DH1 >0 2. Break apart Solvent. • Have to overcome IMFs in the solvent to make room for the solute (expanding the solvent). • DH2 >0

15. 3. Mixing solvent and solute • Allowing the solute and solvent to interact to form the solution (often exothermic) • DH3 depends on what you are mixing. • Molecules can attract each other means DH3 is large and negative. • Molecules can’t attract means DH3 issmall and positive. • This explains the rule “Like dissolves Like”

16. Sep. Solute + Sep. Solvent DH2 Solvent + Sep. Solute DH1 DH3 Solution • Size of DH3 determines whether a solution will form Energy Reactants DHsoln

17. Enthalpy (Heat) of Solution • DHsol’n = DH1 + DH2 + DH3 • DHsol’n may be positive or negative • Enthalpy (heat) of hydration, DHhyd includes steps 2 and 3.

18. Types of Solvent and solutes • If DHsoln is small and positive, a solution will still form because of entropy. • Entropy = measure of disorder • There are many more ways for them to become mixed than there is for them to stay separate.

19. Solution Formation, Spontaneity, and Disorder • Processes in which the energy content of the system decreases tend to occur spontaneously. (tend to be exothermic) • You let go of something, it falls to the floor • Your room gets messy • Change tends to occur in the direction that leads to lower energy or enthalpy for the system. • The amount of disorder in a system a called entropy… processes in which the disorder (entropy) of the system increases tend to occur spontaneously.

20. The solution process involves 2 factors: • Change in enthalpy • Change in entropy • Formation of solutions is favored by the increase in entropy that accompanies mixing… therefore a solution will form unless solute-solute or solvent-solvent IMFs are too strong relative to the solute-solvent interactions.

21. Overview of factors favoring the solution process • Negative DHsol’n • Increase in entropy • For positive values of DHsol’n, it is the increase in entropy that outweighs the increase in energy and causes the solution process to occur.

22. Saturated solutions and Solubility • The process opposite of the solution process is called crystallization. • Solubility = the amount of solute that will dissolve in a solvent at a given temperature. • What is meant by the terms: Saturated, Supersaturated, and Unsaturated

23. Factors Affecting Solubility • Structure (Solute-Solvent Interactions) • Pressure • Temperature

24. Structure (Solute-Solvent Interactions) • “Like dissolves Like” • Polar (hydrophillic) dissolves in polar • Nonpolar (hydrophobic) dissolves in nonpolar • Miscible – two “like” liquids dissolve in one another (Alcohol and Water) • Immiscible – two “un-like” liquids do not dissolve in one another (oil and water).

25. Pressure Effects • Changing the pressure doesn’t effect the amount of solid or liquid that dissolves because they are incompressible. • It does effect gases. • How would an increase in pressure effect gas solubility? • A decrease?

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

27. The gas is at equilibrium with the dissolved gas in this solution. • 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. • This process is known as Henry’s Law.

30. Henry’s Law • the amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution • Sg=kPg • Pg = partial pressure of the gaseous solute above the solution • Sg = concentration of the dissolved gas • k = constant characteristic of a particular solution (dif. for every solution)

31. Sample Henry’s Law problem • Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 of 4.0 atm over the liquid at 25oC. The Henry’s Law constant for CO2 in water at this temperature is 3.1 x 10-2 mol/L-atm.

32. 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.

33. Solubility 100 40 60 80 20 Temperature

34. Temperature Effects • Solubility of gases alwaysdecreases with increasing temperature. • Why do you think that is?

35. Ways of expressing Concentration • Molarity = moles of solute Liters of solution • % mass = Mass of solute x 100 Mass of solution • Mole fraction of component A cA = nA nA + nB + … (Moles of component A divided by total moles of all components)

36. Ways of expressing Concentration • Molality = moles of solute Kilograms of solvent • Molality is abbreviated m • Parts per million (ppm) = mass of component in sol’n x 106 total mass of sol’n

37. Sample problems • A)A solution is made by dissolving 13.5g of glucose (C6H12O6)in 0.100 kg of water. What is the mass percentage of solute in this solution? B) A 2.5g sample of groundwater was found to contain 5.4 mg of Zn+2. What is the concentration of Zn+2 in ppm? • A solution is made by dissolving 4.35g glucose in 25.0 mL of water. Calculate the molality of glucose in the solution.

38. A solution of hydrochloric acid contains 36% HCl by mass. A) Calculate the mole fraction of HCl in the solution. B) calculate the molality of HCl in the solution.

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

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

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

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

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

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

45. Sample problem • Glycerin (C3H8O3) is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25oC. Calculate the vapor pressure at 25oC of a solution made by adding 50.0mL of glycerin to 500.0 mL of water. The vapor pressure of pure water at 25oC is 23.8 torr.

46. Review Question • What is the composition of a pentane-hexane solution that has a vapor pressure of 350 torr at 25ºC ? • The vapor pressures at 25ºC are • pentane 511 torr • hexane 150 torr. • What is the composition of the vapor?

47. Colligative Properties • Because dissolved particles affect vapor pressure, they also affect phase changes. • Colligative Properties are dependent on the number of solute particles, but not their identity. • Boiling-Point elevation • Freezing-Point depression • Osmotic Pressure

48. Boiling-Point Elevation • Since non-volatile solutes lower the vapor pressure, the Boiling Point will rise. Why? • 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

49. Freezing-Point Depression • Because a non-volatile solute lowers the vapor pressure of the solution it lowers the freezing point. Why? • 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

50. 1 atm Vapor Pressure of pure water Vapor Pressure of solution