Unit 6- Solutions

1. Unit 6- Solutions Chapters 14 & 15 Pages 411-469

2. Classification of Matter

3. Focus on Mixtures • Mixtures- Can be separated by physical means into pure compounds or elements • Homogeneous Mixture- • Heterogeneous Mixture- • Classify: salt water, cinnamon and sugar, black coffee, Pepsi.

4. Homogeneous Vs Heterogeneous Mixtures

5. Which is more likely to form homogeneous mixtures? • Mixtures composed of small particles? • Mixtures made up of large particles? When particles of the substances making up a mixture are small, they can be more uniformly mixed, or intermingled. Thus, there is a relationship between particle size and the uniformity of the mixture.

6. Solutions, Suspensions, and Colloidal Dispersions • Solution- homogeneous mixture whose particles are extremely small in size- individual molecules, atoms, or ions. • Suspension- heterogeneous mixture in which one or more of the substances is relatively large. Unless a suspension is stirred or shaken, the large particles sooner or later settle out. • Colloidal Dispersion- heterogeneous mixture in which the particles of one or more of the substances are smaller than those in suspensions, but larger than those in solutions. The particles are large enough to scatter light- Tyndall effect, but too small to filter out.

7. Classification of Matter

8. Focus on Solutions • Solutions are often homogeneous mixtures of two substances. Usually, one of the substances is considered to be dissolved in the other. Solute = Solvent =

9. Consider Gatorade • The solution is a homogenous mixture. • The dissolved particles (sugar, electrolytes) will not come out of solution no matter how long the solution is allowed to stand. • The solution is transparent- the dissolved particles are too small to be seen. • Because of the small size of the dissolved particles, the solution will pass through the finest of filters. • A solution is a homogeneous mixture that is considered to be a single phase even though the compounds may have been in different phases before the solution was formed.

10. Aqueous solutions = water is considered to be the solvent.

11. Types of Solutions • Gas- gas or vapors dissolved in one another- • Air- comprised primarily of O2 and N2 • Liquid solutions- solvent and solution are liquids; solute can be solid, liquid, or gas • vinegar- acetic acid and water, • soft drinks – CO2 and sugar with H2O • Miscible- two liquids that dissolve in each other in all proportions- alcohol and water • Immiscible- liquids that do not dissolve in each other to any degree- water and oil • Solid- two solids uniformly mixed- alloys- Brass- Cu and Zn.

12. Solubility of a Solute

13. Ionic Compounds Dissociate into IonsCovalent (Molecular) Compounds Dissolve

14. Degree of Solubility= How Much Dissolves • “Likes dissolve likes”

15. The effect of pressure on the solubility of a gas

16. Factors Affecting the Rate of Solubility = How Quickly it Dissolves • Rate of solution- a measure of how fast a substance dissolves

17. Solubility and the Nature of a Solvent and a Solute • What holds liquids and solids and to some extent gases together? • What happens when a solid dissolves in a solvent to form a solution? • In order for a solvent to dissolve a solute, the particles of the solvent must be able to separate the particles of the solute and occupy the intervening spaces.

19. Polar Vs Nonpolar Solvents and Solutes • Polar solvents can effectively separate the molecules of other polar substances. • EtOH and Water • NaCl and Water • Positive ions are attracted to negative end of solvent molecules. Negative ions are attracted to the positive ends of solvent molecules. • Dissociation-

22. Nonpolar substances do not dissolve in polar solvents. • Oil and water don’t mix • Butter doesn’t dissolve well in water • Oil/fat have no attraction for polar molecules

23. Soap and Detergents • Surfactants- salts of organic acids having large hydrocarbon like groups. • Soil and dirt are usually held in place by a thin layer of greasy material. Since this layer is nonpolar, the hydrocarbon tails of the detergent ions dissolve in it. The ionic heads of the detergent are hydrated by the water.

25. Energy Changes During Solution Formation • Solid dissolves in a liquid- energy from the heat in the liquid is often required to break the forces of attraction in the solid. • Endothermic = • NH4NO3 dissolving in water. • Most cases in which a solid dissolves in a liquid are endothermic. • Exothermic- • KOH, NaOH, CaCl2 dissolving in water.

26. Solubility Curves and Solubility Tables • What is a major variable we can easily control that affects the solubility of a solid? • Solubility Curve- shows how much solute will dissolve in a given amount of solvent over a range of temperatures. (Table G) • Describe generally what you notice about the solids and their solubilities relative to the temperature of water in the solubility curve. • NaCl, SO2, HCl, NH3

27. Table G: Solubility Curves

28. Concentration of a Solution • Concentration = • Is there a limit to how much Gatorade can be put into water? • Dilute solution- one in which the amount of solute dissolved is small in relation to the amount of solvent. • Concentrated solution- one in which a relatively large amount of solute is dissolved.

29. Saturated – • (REFER TO TABLE G) • Solution equilibrium = the rate at which the undissolved solute goes into solution equals the rate at which the dissolved solute drops out of solution. • Therefore, NO NET CHANGE in amount dissolved. • Unsaturated- • Supersaturated- • UNSTABLE – easily changed to a saturated solution by causing excess solute to precipitate out of solution.

30. Saturated and Unsaturated Vs Dilute and Concentrated KClO3 10oC 5g in 100g of water = NaNO3 10oC 5g in 100g of water = NaNO3 10oC 80g in 100 g of water =? KI 10oC 80 g in 100g water =?

31. Molarity • A method of expressing concentration of a solute in a solvent. Molarity =

32. Molarity 1 M glucose solution = 1 mole of glucose Liter of Solution 1M NaCl = 1 mole Na+ and 1 mole Cl- liter of solution liter of solution

34. Molarity- Example Problems • A chemistry teacher needs to make 500. mL of dilute hydrochloric acid. He puts 0.050 moles of HCl in 500. mL of solution. What is the concentration of the solution? • A household cleaner contains 10.0 g NaOH in a 0.100L solution. What is the molarity of the cleaning solution?

35. Making a Solution • Determine the mass of the solute needed and measure on a balance. • Pour the solute into the proper size volumetric flask. • Add enough distilled water to make the necessary amount of solution. • Cover the volumetric flask and mix completely by inverting.

36. Molarity Example Problem • A student needs to use 0.50 M copper (II) sulfate for a lab. How many grams of copper (II) sulfate pentahydrate must be used to make 250. mL of solution? Briefly describe how the solution would be made.

37. Molarity by Dilution • Less concentrated solutions may be prepared from concentrated solutions using a dilution formula. Where M1 describes the molarity of the stock solution V1 is the volume of the stock solution M2 and V2 describe the molarity and volume of the more dilute solution you’re making

38. Making a Dilution • 1. Determine the volume of concentrated solution needed and measure in a graduated cylinder or volumetric pipet. • 2. Pour the more concentrated solution into the proper size volumetric flask. • 3. Add enough distilled water to make the necessary amount of solution. • 4. Cover the volumetric flask and mix completely by inverting.

40. Dilution Problem Example • 0.10M HCl is needed for a lab. How would 250. mL of 0.10 M HCl be prepared from 6.0 M HCl stock solution?

41. Molality • Molality (m) = the number of moles of solute dissolved in each kilogram of solvent.

42. Example Molality Problem • Calculate the molality of a solution prepared by dissolving 58.44 g of NaCl in 500. g of water.

43. Quick Review of Electrolytes Vs Non Electrolytes

44. Colligative Properties • A property that depends on the amount of solute particles but is independent of their identity. In other words, how much, not what it is. • Vapor Pressure- solute molecules take up space, preventing some molecules from leaving the liquid. • Osmotic Pressure- pressure required to prevent osmosis. Adding a solute to a solvent increases the osmotic pressure of a solution.

45. Vapor Pressure Osmotic Pressure

46. Why it’s a bad idea to drink seawater!

47. 3. Boiling Point Elevation- related to vapor pressure reduction. Higher temperatures are necessary to get solvent molecules up into the gas phase. When vapor pressure equals atmospheric pressure, the substance boils. If vapor pressure is reduced by the addition of a solute, more energy is required to get the substance to boil. 4. Freezing point depression- similar to boiling point elevation. Solute molecules prevent the formation of intermolecular forces between solvent molecules forming the solid phase.

48. The Effect of a Solute on Vapor Pressure

49. Boiling Point Elevation • The amount by which the temperature of a solution increases as a result of addition of a solute is called “boiling point elevation” or DTb • DTb is directly proportional to molality Where Kb = molal boiling point constant m = molality of the solution d.f. = dissociation factor (electrolytes only)

50. Example Boiling Point Elevation • How many grams of benzoic acid (C7H6O2) must be dissolved in 79.1 g of ethanol to raise the boiling point by 4.00oC? Kb for ethanol is 1.20oC/m.