1. I. The Nature of Solutions Solutions

2. Mixtures Review  Homogeneous –Solutions  Heterogeneous –Suspension - is a fluid containing solid particles that are sufficiently large enough to settle. –Colloid - is a substance microscopically dispersed evenly throughout another substance. Colloid particles are large enough to scatter light, this is called the Tyndall effect. The Tyndall effect is used to distinguish between a colloid and a solution. –Emulsion - is a mixture of two or more immiscible (un-blendable) liquids.

3. Solutions  Solution - homogeneous mixture Solvent - present in greater amount Solute - substance being dissolved

4. Solutions Solute - KMnO 4 Solvent - H 2 O

5. Solvation  Solvation – the process of dissolving solute particles are separated and pulled into solution solute particles are surrounded by solvent particles

6. Solvation  Dissociation –separation of an ionic solid into aqueous ions NaCl(s)  Na + (aq) + Cl – (aq)

7. Solvation  Ionization –breaking apart of some polar molecules into aqueous ions HNO 3 (aq) + H 2 O(l)  H 3 O + (aq) + NO 3 – (aq)

8. B. Solvation  Molecular Solvation –molecules stay intact C 6 H 12 O 6 (s)  C 6 H 12 O 6 (aq)

9. B. Solvation Strong Electrolyte Non- Electrolyte solute exists as ions only - + salt - + sugar solute exists as molecules only - + acetic acid Weak Electrolyte solute exists as ions and molecules DISSOCIATIONIONIZATION View animation online.animation

10. Solubility  Soluble – substance can dissolve in a solvent Ex: salt in water  Insoluble – substance cannot dissolve in a solvent Ex: sand in water Ex: sand in water

11. “like dissolves like”  rule used to determine if substance will dissolve in another  based on attractive forces between solute and solvent

12. “like dissolves like”  polar solvents – dissolve polar molecular compounds and ionic compounds ex: salt and water, alcohol and vinegar  nonpolar solvents – dissolve nonpolar compounds ex: oil and gasoline

13. B. Solvation NONPOLAR POLAR “Like Dissolves Like”

14. Factors Affecting Rate of Solvation  How can you dissolve something faster? a) increase temp of solvent this accelerates particles creating more particle collisions The higher the temperature, higher the solubility (for most cases)

15. Factors Affecting Rate of Solvation b) agitate the solution more particle collisions between solute and solvent more particle collisions between solute and solvent c) Increase surface area of solute breaking into smaller pieces allows more solute to be in contact with solvent

16. Heats of Solution  Formation of a solution is accompanied by an energy change.  The formation of a liquid-solid solution can absorb or release heat.  Solvent molecules are held together by IMF, solvent-solvent attraction.  Solute molecules are also held together by IMF, solute-solute attraction.

17. Heats of Solution  Energy is required to separate these molecules in order for solvation to occur.  Heat of Solution is the net amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent.  Heat of solution is negative when heat is released and positive when heat is absorbed.

18. Solubility  maximum amount of solute that can dissolve in a solvent at a specific temperature how much solute can be put into solvent?

19. Unsaturated Solution  less than maximum amount of solute dissolved Ex. if I put sugar into water and all sugar is dissolved, solution is unsaturated

20. Saturated Solution  contains maximum amount of solute dissolved Ex: if I put sugar into water and not dissolves (you can see the sugar), the solution is saturated

21. Supersaturated Solution  contains more solute than a saturated solution at the same conditions Ex: a saturated solution made at high temp cools slowly. Slow cooling allows excess solute to remain dissolved in solution at lower temperature very unstable

22. Solubility Solubility SATURATED SOLUTION no more solute dissolves UNSATURATED SOLUTION more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form concentration

23. Solubility  Solubility –maximum grams of solute that will dissolve in 100 g of solvent at a given temperature –varies with temp –based on a saturated solution

24. Solubility  Solubility Curve –shows the dependence of solubility on temperature

25. Solubility Curve Saturated- Line represents max amount solute that will dissolve at a given temperature Temperature Solubility (g solute/ 100 g H 2 O) Unsaturated (below line) Supersaturated (above line)

26. Solubility  Solids are more soluble at... –high temperatures.  Gases are more soluble at... –low temperatures & –high pressures (Henry’s Law). –EX: soda

27. Henry’s Law  Henry's Law states that the mass of a gas which will dissolve into a solution is directly proportional to the partial pressure of that gas above the solution.  Applies to gas-liquid solutions at constant temperature

28. Solubility Rules  Solubility is a result of an interaction between polar water molecules and the ions which make up a crystal.  These two forces determine the extent to which solution will occur.

29. Solubility Rules  It is difficult to quantify how these two forces will work together.  Therefore, we refer to a set of generalizations about solubility  These are called 'solubility rules',  These rules are based upon experimentation.

30. Solubility Rules  NO 3 - - All nitrates are soluble.  Cl - - All chlorides are soluble except AgCl, Hg 2 Cl 2, and PbCl 2.  (SO 4 ) 2 - - Most sulfates are soluble. Exceptions include BaSO 4, PbSO 4, and SrSO 4.  (CO 3 ) 2 - - All carbonates are insoluble except NH 4 + and those of the Group 1 elements.  OH - - All hydroxides are insoluble except those of the Group 1 elements, Ba(OH) 2, and Sr(OH) 2. Ca(OH) 2 is slightly soluble.  S 2 - - All sulfides are insoluble except those of the Group 1 and Group 2 elements and NH 4 +.

31. II. Concentration Solutions

32. Solution Concentration  Concentration – how much solute dissolved in amount of solvent what is difference between concentrated and diluted?

33. Concentrated vs. Dilute

34. What is different between the glasses of Kool-aid?

35. Solution concentration can be described generally  Dilute - reduced in strength, weak, watered down.  Concentrated – stronger, pure. Has less water.

36. What is the problem with just using dilute and concentrated as descriptions of the solution concentration?

37. Is solution B dilute or concentrated?  The terms dilute and concentrated are relative.  Scientists need a more precise way of referring to the concentration of a solution. ConcentratedDilute Solution A Solution B Solution C

38. Concentration  The amount of solute in a solution.  Describing Concentration –% by mass - medicated creams –% by volume - rubbing alcohol –ppm, ppb - water contaminants –molarity - used by chemists –molality - used by chemists

39. Molarity The ratio of the moles of solute to the volume of solvent in liters. Molarity (M) = Moles of solute Volume in Liters of Solvent

40. How to read Molarity  6.0 M NaCl  Read: “6 molar solution of NaCl”  Can be abbreviated 6M solution  You must be careful to label the molarity with a capital M so that it is not confused with m for molality.

41. How to make a solution using molarity (6M NaCl) (a) Add 6 moles NaCl to the volumetric flask. How would you measure that? 6moles NaCl 58.443 g NaCl 1 mole NaCl = 351 g NaCl

42. How to make a solution using molarity (6M NaCl) (b) Add dH 2 O to dissolve and mix the NaCl (c) Fill the flask with dH 2 O until you reach the 1000mL line.

43. Types of Calculations with Molarity: 1. Finding concentration of a solution. 2. Finding the mass of solute needed. 3. Finding the volume of solution made.

44. Finding Concentration  Antifreeze is a solution of ethylene glycol, C 2 H 6 O 2 in water. If 4.50 L of antifreeze contains 27.5 g of ethylene glycol, what is the concentration of the solution? 27.5 g C 2 H 6 O 2 62.08 g C 2 H 6 O 2 1 mol C 2 H 6 O 2 4.5 L = 0.0984 mol/L or 0.0984 M C 2 H 6 O 2

45. Finding Mass  What mass of sodium carbonate, Na 2 CO 3, is present in 50 ml of a 0.750M solution? 50 ml 1 L 1000 mL 0.750 mol 1 L 1 mol Na 2 CO 3 105.99 g Na 2 CO 3 = 3.97 g Na 2 CO 3 Conversion Factor

46. Finding Volume  What volume of 1.50 mol/L HCl solution contains 10.0 g of hydrogen chloride? 10.0 g HCl 1 mol HCl 36.46 g HCl 1.50 mol 1 L = 0.183 L or 183 mL Conversion Factor

47. Molarity Practice Problems

48. Practice Problems 1. A 0.750 L aqueous solution contains 90.0 g of ethanol, C 2 H 5 OH. Calculate the molar concentration of the solution in mol·L -1.

49. Practice Problem 2. What mass of NaCl are dissolved in 152 mL of a solution if the concentration of the solution is 0.364 M?

50. Practice Problem 3. What mass of dextrose, C 6 H 12 O 6 is dissolved in 325 mL of 0.258 M solution?

51. Practice Problem 4. A mass of 98 g of sulfuric acid, H 2 SO 4, is dissolved in water to prepare a 0.500 M solution. What is the volume of the solution?

52. Practice Problem 5. A solution of sodium carbonate, Na 2 CO 3, contains 53.0 g of solute in 215 mL of solution. What is its molarity?

53. Practice Problem 6. What is the molarity of a solution of HNO3 that contains 12.6 g of solute in 5.00 L of solution?

54. Molality mass of solvent only 1 kg water = 1 L water

55. Molality  Find the molality of a solution containing 75 g of MgCl 2 in 250 mL of water. 75 g MgCl 2 1 mol MgCl 2 95.21 g MgCl 2 = 3.2 m MgCl 2 0.25 kg water

56. Molality  How many grams of NaCl are req’d to make a 1.54m solution using 0.500 kg of water? 0.500 kg water 1.54 mol NaCl 1 kg water = 45.0 g NaCl 58.44 g NaCl 1 mol NaCl

57. Dilution  Preparation of a desired solution by adding water to a concentrate.  Moles of solute remain the same.

58. Dilution  What volume of 15.8M HNO 3 is required to make 250 mL of a 6.0M solution? GIVEN: M 1 = 15.8M V 1 = ? M 2 = 6.0M V 2 = 250 mL WORK: M 1 V 1 = M 2 V 2 (15.8M) V 1 = (6.0M)(250mL) V 1 = 95 mL of 15.8M HNO 3

59. Preparing Solutions  500 mL of 1.54M NaCl 500 mL water 45.0 g NaCl –mass 45.0 g of NaCl –add water until total volume is 500 mL –mass 45.0 g of NaCl –add 0.500 kg of water 500 mL mark 500 mL volumetric flask  1.54m NaCl in 0.500 kg of water

60. Preparing Solutions  250 mL of 6.0M HNO 3 by dilution –measure 95 mL of 15.8M HNO 3 95 mL of 15.8M HNO 3 water for safety 250 mL mark –combine with water until total volume is 250 mL –Safety: “Do as you oughtta, add the acid to the watta!”

61. Colligative Properties

62. Colligative Properties of Solutions  Solutes affect the physical properties of their solvents  Colligative properties – properties that depend only on the number of solute particles present, not their identity  Ex: boiling point, freezing point

63. Electrolytes  Definition – substances that break up (ionize) in water to produce ions; can conduct electricity - consist of acids, bases, ionic compounds Ex: NaCl  Na 1+ + Cl 1- H 2 SO 4  2 H + + SO 4 2-

64. Nonelectrolytes  Definition – do not break up (ionize) in water, they stay the same; do not conduct electricity - usually molecular/covalent compounds Ex:sugar C 6 H 12 O 6  C 6 H 12 O 6 ethanolC 2 H 5 OH  C 2 H 5 OH

65. Properties of Water

66. Water is Polar Water is polar because oxygen is a “bully” and does not share the electrons. Water is polar because oxygen is a “bully” and does not share the electrons. Oxygen is negative with more electrons and Hydrogen is positive with less electrons. Oxygen is negative with more electrons and Hydrogen is positive with less electrons.

67. Cohesion  Cohesion is the attraction of the same type of molecules to each other.  Water molecules stick together and “hold hands” with hydrogen bonds

68. Water is the Universal Solvent Water is polar and can dissolve both salts, sugars, and other molecules.

69. Adhesion  Adhesion is the attraction of molecules between two different substances.  Example: The attraction of water molecules to a glass.

70. Adhesion Adhesion causes the water molecules to stick to the sides of the graduated cylinder creating a meniscus. Adhesion causes water to rise in a straw.