1. Chemistry Chapter 12 - Solutions Right now: Make a list of solutions 1

2. Chapter 12 Assignments 12.1 Review 103/1-6 12.2 Review 105/1-4 12.3 Review 107/1-5 110/4-6 2

3. Types of Mixtures Solution – homogeneous mixture of two or more substances in a single phase. Solvent – the substance in which the solute dissolves. Solute – the substance, of which there is less, that dissolves in the solvent. Suspensions – large particles in a solvent where they settle out (because of gravity) unless continually mixed. Colloids – mixtures where particles are larger than those in a solution but smaller than in a suspension. 3

4. Table 1 - Types of solutions Solute StateSolvent stateExample gas oxygen in nitrogen gasliquidCO 2 in water liquid alcohol in water liquidsolidHg in Ag and Sn (dental amalgam) solidliquidsugar in water solid Cu in Ni (alloy) 4

5. Table 2 – Classes of Colloids Class of colloidPhasesExample Solsolid dispersed in liquidpaints, mud Gelsolid network extending throughout liquid gelatin Liquid emulsionliquid dispersed in a liquidmilk, mayonnaise Foamgas dispersed in liquidsharing cream, whipped cream Solid Aerosolsolid dispersed in gassmoke, auto exhaust Liquid Aerosolliquid dispersed in gasfog, mist, clouds, aerosol spray Solid Emulsionliquid dispersed in solidcheese, butter 5

6. Testing for Colloids – Tyndal Effect A beam of light will pass through a solution A beam of light passed through a colloid will scatter the light, making the beam visible. http://pixshark.com/tyndall-effect-fog.htm 6

7. Conduct Electricity? Electrolytes - Those solutions that will conduct electricity. These are usually formed when ionic compounds (such as NaCl) or when highly polar molecules (such as HCl) are dissolved in water. Nonelectrolytes – those that will NOT conduct electricity. A sample is sugar (a nonpolar colvalent substance) dissolved in water. 7

8. Thursday, 3.26.2015 Extra Credit (3 points) – Draw a science related picture/figure/? To enter in our Fine Arts Festival. Due by Monday, 3/30. Lab 5 due tomorrow ASOP. Perform Lab 6 tomorrow. No prelab. Yeah!!! Remember to bring 3 shiny pennies or else you will have to scratch off oxidation with steel wool. Needed today: pencil, 1 sheet paper Have Ch 12.1 review page on your desk and quickly show me that you did it (both sides). 8

9. Factors Affecting Dissolving Temperature Surface Area Stirring 9

10. Temperature has an effect on solubility 210 g sugar 100 mL H 2 O 20 o C Undis- solved sugar 210 g sugar 100 mL H 2 O 30 o C All the sugar is dis- solved 10 solubility: the amount of a solute that will dissolve in a particular solvent at a particular temperature and pressure.

11. Temperature and solubility Temperature does not have the same effect on the solubility of all solutes 11

12. The rate of solubility increases: - with an increase in temperature At higher temperatures: - solid solutes (like salt and sugar) are more soluble - gases are less soluble with an increase in surface area of the solute 12

13. Seltzer water is a supersaturated solution of CO 2 in water supersaturation: term used to describe when a solution contains more dissolved solute than it can hold. This solution is unstable, and the gas “undissolves” rapidly (bubbles escaping) 13

14. Solubility Values The solubility of a substance is the amount of that substance required to form a saturated solution with a specific amount of solvent at a specified temperature. 14 Substance020406080100 C 12 H 22 O 11 179204238287362487 Li 2 CO 3 1.541.331.171.010.850.72 Table 4 Solubility of Solutes as a Function of Temperature in g solute/100. g H 2 0.

15. Like Dissolves Like Dissolving depends on: Bonding type Polar or nonpolar Intermolecular forces between solute and solvent. Ionic compounds are generally soluble in water and the ions become hydrated. Ionic compounds are generally NOT soluble in nonpolar solvents. 15

16. In general, “like” dissolves “like” Polar solvents dissolve polar solutes Nonpolar solvents dissolve nonpolar solutes Not everything dissolves in water. Why not? 16

17. Solvents Immiscible – liquids that are not soluble in each other. Miscible – liquids that dissolve freely in one another in any proportion. What solvents do you have at your house? Mineral oil, turpentine, paint remover, nail polish remover Water is called the universal solvent because it dissolves both ionic and covalent compounds. 17

18. Two special solvents Ethanol and Toluene http://www.easychem.com.au/production-of- materials/renewable-ethanol/ethanol-as-a- solvent http://www.easychem.com.au/production-of- materials/renewable-ethanol/ethanol-as-a- solvent http://toxtown.nlm.nih.gov/text_version/che micals.php?id=30 http://toxtown.nlm.nih.gov/text_version/che micals.php?id=30 18

19. For each of these do the following on paper. Chemical Name Chemical Formula Lewis Diagram (show lone pairs) Structural diagram (3 D) Polar or Nonpolar and why Uses Concerns (due eop) 19

20. 20 Ethanol Structural Formula – show lone pairs on this  Structural formula 3D

21. 21 Toluene

22. Water as a solvent hydration: the process of molecules with any charge separation to collect water molecules around them. Not chemically bonded 22

23. Reactions in liquids Life involves many complex chemical reactions that only occur in aqueous solutions! A step in the Krebs cycle – this is how energy is extracted from glucose 23

24. Pressure and solubility The solubility of gases are affected greatly by pressure (not solids and liquids). When a soda bottle is opened, the pressure is decreased and the gas (CO 2 ) can escape from the liquid. That is the “fizzing”. Henry’s Law – The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid. 24

25. Enthalpy of Solution The amount of energy released or absorbed at heat when a specific amount of solute dissolves in a solvent. 25 http://blog.science-matters.org http://openstudy.com

26. concentration: the amount of each solute compared to the total solution. 12.3 Concentration & Solubility 26

27. There are several ways to express concentration 27 molality, m = moles solute kg solvent

28. Suppose you dissolve 10.0 g of sugar in 90.0 g of water. What is the mass percent concentration of sugar in the solution? Asked:The mass percent concentration Given:10 g of solute (sugar) and 90 g of solvent (water) Relationships: Solve: 28

29. Calculate the molarity of a salt solution made by adding 6.0 g of NaCl to 100 mL of distilled water. Asked:Molarity of solution Given:Volume of solvent = 100.0 mL, mass of solute (NaCl) = 6.0 g Relationships: Formula mass NaCl Form. Mass = 22.99 + 35.45 = 58.44 g/mole 1,000 mL = 1.0 L, therefore 100 ml = 0.10 L M = moles = 6.0g NaCl 1 mole NaCl = 1.03 M NaCl L LL LlLlLL0.100 L llll58.44 g NaCl 29

30. 30 Calculate the molality, m of a solution containing 350.9 g of NaCl in 750.0 g of water. (Density of water is 1g/mL.) Find the molar mass of NaCl: 22.99 + 35.45 = 58.44 g/mol Mass to mole conversion: 350.9 g NaCl 1 mole = 6.004 mole NaCl 58.44g m = moles solute = 6.004 mole NaCl = 8.005 m kg solvent 0.7500kg

31. 100 mL H 2 O What happens when you add 10 g of sugar to 100 mL of water? Conc. (%) = 10 g/110 g Water molecules dissolve sugar molecules sugar 10 g 31

32. What happens when you add 10 g of sugar to 100 mL of water? But when two sugar molecules find each other, they will become “undissolved” (solid) again… … then, they become redissolved in water again. 32

33. What happens when you add 10 g of sugar to 100 mL of water? This is an aqueous equilibrium! 33 saturation: situation that occurs when the amount of dissolved solute in a solution gets high enough that the rate of “undissolving” matches the rate of dissolving.

34. Preparing a solution How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. M, molarity = moles solute / liter solution 1.Determine the formula mass of the solute. Molar mass of CaCl 2 34

35. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. Molar mass of CaCl 2 : 110.98 g/mole How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. We need 0.5 moles CaCl 2 35

36. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. Molar mass of CaCl 2 : 110.98 g/mole How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. We need 0.5 moles CaCl 2 We need 55.49 g CaCl 2 36

37. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. 3.Weigh the grams of solute on the balance. How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. 37

38. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. 3.Weigh the grams of solute on the balance. 4.Add the solute to a volumetric flask or graduated cylinder. How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. 38

39. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. 3.Weigh the grams of solute on the balance. 4.Add the solute to a volumetric flask or graduated cylinder. 5.Fill the flask about two thirds of the way up with distilled water. How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. Do not fill all the way up 500.0 mL mark 39

40. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. 3.Weigh the grams of solute on the balance. 4.Add the solute to a volumetric flask or graduated cylinder. 5.Fill the flask about two thirds of the way up with distilled water. 6.Mix the solution until the solid dissolves completely. How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. 40

41. Preparing a solution 1.Determine the formula mass of the solute. 2.Use the formula mass of the solute to determine the grams of solute needed. 3.Weigh the grams of solute on the balance. 4.Add the solute to a volumetric flask or graduated cylinder. 5.Fill the flask about two thirds of the way up with distilled water. 6.Mix the solution until the solid dissolves completely. 7.Fill the volumetric flask or graduated cylinder up to the correct volume marker. How to prepare a 500.0 mL solution of a 1.0 M CaCl 2 solution. 41

42. Ways to express concentration: A higher temperature causes higher: - solubility of solutes how much - rates of solubility how fast 42 Molality, m = moles solute kg solvent

43. 43 What do you remember? M stands for molarity m stands for molality Molarity is moles solute/L solution Molality is moles solute/kg solvent

44. 44 Two other ways to determine concentration: g/L % when comparing mass of solvent with mass of solution Methods to dissolve faster: stirring, heating, smaller pieces (yielding larger surface area)

45. Cleo, Will you go to the prom with me? Yours forever, Tony 45

46. Antonio, I thought you would never ask! Cleopatra 46

47. Chapter 13 Ions in Aqueous Solutions & Colligative Properties Book Asgns: 458/1,2(a-f),3,5,8 459/18-20,22 459/26,28,37,39,41 47

48. Compounds in Aqueous Solutions Reaction rate is generally dependent upon concentration – greater concentration means reaction occurs faster Heat of solution – energy absorbed or released when a solute dissolves in a particular solvent exothermic, loss of energy (gives off energy) or negative heat of solution (feels hot) endothermic, energy absorbed (feels cold) or positive heat of solution 48

49. 49 Exothermic – energy lost Endothermic – energy gained

50. NH 4 NO 3 (s) + H 2 O(l) → NH 4 + (aq) + NO 3 – (aq)∆H = +25.7 kJ/mole HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l)∆H = –56 kJ/mole What changes is the enthalpy enthalpy: the energy potential of a chemical reaction measured in joule per mole (J/mole) or kilojoules per mole (kJ/mole). 50

51. ∆H reaction = –56 kJ/mole ∆H solution = +56 kJ/mole Opposite signs! Heat released by the reaction = Heat gained by the solution HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l) ∆H = –56 kJ/mole 51

52. Volumes of solute and solvent do not add up to the volume of solution 20 g salt80 mL water 87 mL solution! Salt dissociates into ions, which fit in between water molecules Solution vs. pure solvent 52

53. 13.1 Compounds in Aqueous Solutions Dissociation is the separation of ions that occurs when an ionic compound dissolves. Ionization occurs when covalent compounds dissolve in a solvent. 53

54. Electrolyte solutions Aqueous solutions containing dissolved ions are able to conduct electricity 1 mole of solute → 2 moles of ions 1 mole of solute → 3 moles of ions The greater the number of particles in solution, the greater the effects. 54

55. 55 Table 1 General Solubility Guidelines 1. Sodium, potassium, and ammonium compounds are soluble in water. 2. Nitrates, acetates, and chlorate are soluble. 3. Most chlorides are soluble, except those of silver, mercury (I), and lead. Lead (II) chloride is soluble in hot water. 4. Most sulfates are soluble, except those of barium, strontium, lead, calcium, and mercury. 5. Most carbonates, phosphates, and silicates are insoluble, except those of sodium, potassium, and ammonium. 6. Most sulfides are insoluble, except those of calcium, strontium, sodium, potassium, and ammonium.

56. 436 Practice 56

57. Net Ionic Equations Identify the precipitate that forms when aqueous solutions of potassium sulfate and barium nitrate are combined. Write the net ionic equation. K 2 SO 4 (aq) +Ba(NO 3 ) 2 (aq) --> KNO 3 (aq)+BaSO 4 (s) 2K +1 (aq) + SO 4 –2 (aq) + Ba +2 (aq) + 2NO 3 -1 (aq) --> 2K +1 (aq) + 2NO 3 -1 (aq) +BaSO 4 (s) Net ionic equation: SO 4 –2 (aq) + Ba +2 (aq) --> + BaSO 4 (s) 57

58. 440 - Net Ionic Equations Practice. Solubility Chart needed. 1. Will a precipitate form if solutions of potassium sulfate and barium nitrate are combined? If so, write the net ionic equation for the reaction. 3. Will a precipitate form if solutions of barium chloride and magnesium sulfate are combined? If so, write the net ionic equation for the reaction. 58

59. Why does ice melt when salt is sprinkled on it? 59 13.2 Colligative Properties of Solutions

60. Freezing point depression Why does ice melt when salt is sprinkled on it? Pure water freezes at 0 o C, but a water and salt solution freezes at a lower temperature. 60

61. colligative property: physical property of a solution that depends only on the number of dissolved solute particles not on the type (or nature) of the particle itself. ** Pure solvent Solid formation is not hindered Solution Solute particles “get in the way” of solid formation OrderEntropy more lessmore less 61 ** Affects on the solvent with a dissolved solute: Boiling point will elevate (  t b ) and the Freezing point will depress (  t f ).

62. To calculate the freezing point of a solution: Do not get confused with molarity, M (moles solute / L of solution) 62 Freezing point depression constant  T f = K f x m Change in freezing molality Point, o C

63. Calculate the freezing point of a 1.8 m aqueous solution of antifreeze that contains ethylene glycol (C 2 H 6 O 2 ) as the solute. Asked: The freezing point of a 1.8 m solution of ethylene glycol Given: molality, m = 1.8 m; K f = 1.86 o C/m (K f, freezing point depression for water, the solvent – see chart on pg 448.) The freezing point is lowered by 3.35 o C. 63

64. 64 Freezing Point Depression Problem: Use table 2, page 448. A solution consists of 10.3 g of the nonelectrolyte glucose, C 6 H 12 0 6, dissolved in 250. g water. What is the freezing point depression of the solution? Mass = 10.3 g C 6 H 12 0 6 --> moles 250. g H 2 0 --> kg = 0.250 kg H 2 0 Find m --> find  t f m = mol solute = 10.3 g C 6 H 12 0 6 1 mol = 0.229 m kg solvent 0.250 kg H 2 0 180 g  t f = m K f = 0.229 m -1.86 o C = - 0.426 o C m Molar mass of C 6 H 12 0 6 6C6x12 = 72 12 H 12x1 = 12 6 (0) 6x16 = 96 180 g/mol

65. 65 Freezing Point Depression Problem: Use table 2, page 448. 2. In a laboratory experiment, the freezing point of an aqueous solution of glucose is found to be -0.325 o C. What is the molal concentration of this solution?

66. 66 Freezing Point Depression Problem: Use table 2, page 448. 3. If 0.500 mol of a nonelectrolyte solute are dissolved in 500.0 g of ether, what is the freezing point of the solution? 0.500 mol solute 0.500 kg ether 0.5 / 0.5 = 1 molal  t f =? o C K f ether = -l.79 o C / m  t f = m K f = (1 m )(-l.79 o C / m) = -l.79 o C New freezing pt = (-116.3 – 1.79) o C = -118.1 o C

67. Ionization occurs when covalent compounds dissolve in a solvent. Water reacting with molecular compounds containing H can release a H + (a proton) and form an H 3 0 + ion (hydronium ion). Sample: HCl + H 2 0 --> H 3 0 + + Cl - The HCl is ionizing. 67

68. Strength of Electrolytes Strong Electrolytes dissolve well in aqueous solution and conduct electricity well even when in dilute solution. (HCl, HBr, HI) HCl + H 2 0 --> Cl - + H 3 0 + Weak Electrolytes don’t dissolve well in aq. soln., conduct electricity poorly, and contain ions and molecules.(HF, HC 2 H 3 0 2 ) HF + H 2 0 H 3 0 + + F – HC 2 H 3 0 2 + H 2 0 H 3 0 + + C 2 H 3 0 2 - 68

69. Ch 13 Assignments From the book: 458/1,2(a-f),3,5,8 459/18-20,22 459/26,28,37,39,41 69

70. Reaction rates increase with: increasing concentrations increasing temperatures Solution vs. pure solvent density (solution) > density (pure solvent) colligative properties: freezing point depression is an example 70

71. 71 molality, m = moles solute kg solvent Ch 12 & 13 Formula list continued: