Unit 11- Solubility & Solutions, Ch. 17 & 18

1. Unit 11- Solubility & Solutions, Ch. 17 & 18

2. I. Water A. The Molecule 1. O—H bond is highly polar 2. Bond angle 105° making it Bent shaped 3. Water Molecule as a whole is polar 4. Attracted to each other by intermolecular hydrogen bonds Greater electronegativity

3. I. Water (cont.) B. Important Properties 1. High surface tension 2. low vapor pressure • hydrogen bonds hold molecules to one another, tendency to escape surface is low 3. high specific heat capacity • 4.184 J/g×°C 4. high melting and boiling points • 0°C and 100°C

4. I. Water (cont.) C. Surface Tension– inward force, or pull, that tends to minimize the surface area of a liquid • Surfactant – wetting agent such as soap or detergent that decreases the surface tension by interfering with hydrogen-bonding Responsible for high surface tension

5. I. Water (cont.) D. Atypical Ice 1. As a typical liquid cools, density increases b/c Volume decreases as the mass stays constant 2. As water cools it first behaves like a typical liquid until it reaches 4°C 3. Below 4°C the density of water starts to decrease **Ice is one of only a few solids that float in their own liquid.

6. Atypical Ice Do not need to write. °

7. Atypical Ice Why does ice behave so differently? Open framework arranged like a honeycomb. Framework collapses, molecules packed closer together, making it more dense

8. II. The Solution Process A. Solution - homogeneous mixture consisting of the same properties throughout (exists in a single phase) Solute - substance being dissolved - smaller quantity substance Solvent - dissolving medium - larger quantity substance - usually water

9. II. The Solution Process (cont.) • Aqueous Solution (aq) – a solution in which the solvent is water HCl(g) + H2O(l) → H3O(aq) + Cl-(aq) potassium chromate, K2CrO4 (yellow) cobalt(II) nitrate, Co(NO3)2 (red) copper(II) sulfate, CuSO4 (blue) potassium permanganate, KMnO4 (purple) potassium dichromate K2Cr2O7 (orange) nickel(II) chloride, NiCl2 (green)

10. II. The Solution Process (cont.) B. Solvation– the process of dissolving 1st solute particles (salt) are surrounded by solvent particles (water) First... 2nd solute particles (salt) are separated and pulled into solution (salt water) Then...

11. NONPOLAR NONPOLAR POLAR POLAR II. The Solution Process (cont.) C. “Like Dissolves Like” Polar solvents dissolve polar molecules and ionic compounds Nonpolar solvents dissolve nonpolar compounds

12. “Polar Dissolves Polar”

13. Polar vs. Nonpolar • Polar Molecule – uneven electron forces acting on the central atom. • Water • Nonpolar Covalent Bond – when 2 atoms are joined by a covalent bond and the bonding electrons are shared equally

14. III. Water of Hydration • Water molecules that are integral part of the crystal structure • Compound that contains water of hydration is called a hydrate • (ex. CuSO4.5H2O) • Effloresce – losing the water of hydration • Deliquescent– removes water from air to form a solution

15. III. Water of Hydration • Calculate the percent by mass of water in sodium carbonate decahydrate (Na2CO3·10H2O) Percent H2O = mass of water x 100% mass of hydrate Mass of 10H2O = 180 g Mass of Na2CO3·10H2O = 286g Percent H2O = 180 g x 100% = 62.9% or 286 g 63%

16. IV. Electrolytes A. Electrolytes – compounds that conduct an electric current in solutions • All ionic compounds are electrolytes • Compounds that don’t conduct an electric current are called nonelectrolytes – not composed of ions, includes many molecular compounds (covalent compounds)

17. IV. Electrolytes (cont.) - + - - + + water sugar salt Weak Electrolyte Non- Electrolyte Strong Electrolyte solute exists as ions and molecules solute exists as molecules only solute exists as ions only

19. V. Heterogeneous Mixtures A. Suspensions – mixtures from which particles settle out upon standing and the average particle size is greater than 100 nm in diameter. Clearly identified as two substances Gravity or filtration will separate the particles B. Colloids – heterogeneous mixtures containing particles that are between 1 nm and 100 nm in diameter Appear to be homogeneous but particles are dispersed through medium Ex: paint, aerosol spray, smoke, marshmallow, whipped cream

20. V. Heterogeneous Mixtures C. Tyndall Effect - phenomenon observed when beam of light passes through a colloid or suspension Colloids exhibit the Tyndall effect Colloid Solution

21. VI. Solubility • defined as the maximum grams of solute that will dissolve in 100 g of solvent at a given temperature • based on a saturated solution

22. Solubility SATURATED SOLUTION max amount no more solute dissolves SUPERSATURATED SOLUTION over max amount becomes unstable, crystals form UNSATURATED SOLUTION capable of dissolving more solute • Supersaturated solutions are not in equilibrium with the solid substance and can quickly release the dissolved solids. • Saturated solution is one that is in equilibrium with respect to the dissolved substance. These conditions can quickly change with temperature. Concentration Increasing

23. VI. Solubility (cont.) B. Factors Affecting Solubility 1. Stirring (agitation) • Increases solubility b/c fresh solvent is brought in contact with the surface of the solute • Surface phenomenon

24. VI. Solubility (cont.) 2. Temperature • Increases solubility by increasing kinetic energy, which increases the collisions b/w molecules of solvent and the surface of the solute • Increases amount and rate of solute dissolved

25. VI. Solubility (cont.) 3. Surface Area • A smaller particle size dissolves more rapidly than larger particles size • Surface phenomenon • More surface area exposed, faster rate of dissolving

26. VI. Solubility (cont.) A. Solubility Curve • shows the dependence of solubility on temperature • Note: the solubility of the gases are greater in cold water than in hot water Out of the solids, which has the lowest solubility at 40°C? KClO3 Which has the highest solubility at 20°C? KI

27. VI. Solubility (cont.) How many grams of KNO3 can be dissolved in 100g of water at 60°C? 110g What is the solubility of NH4Cl at 50°C? 50 g/L Which salt shows the least change in solubility from 0°C to 100°C? NaCl

28. Output side 1. Which of the following compounds dissolved the highest at 20°C? 2. The lowest at 20°C? 3. Overall which compound dissolved the fastest from 0°C to 100°C? 4. Name a compound in the graph that is a gas? How do you know it’s a gas?

29. VII. Concentrations of Solutions • Concentration of a solution is a measure of the amount of solute that is dissolved in a given quantity of solution. • Dilute solution – contains a low concentration of solute. • Concentrated solution – contains a high concentration of solute. • Molarity (M) – number of moles of a solute dissolved per liter of solution • a.k.a. molar concentration

30. A. Molarity • Calculates the number of moles in 1 L of the solution Molarity (M) = moles of solute liters of solution Ex 1 Calculate the molarity when 2 mol of glucose is dissolved in 5 L of solution. 2 mol glucose 5 L solution M= n L M= n L = 0.4 mol/L = 0.4 M

31. Ex 2 How many moles of solute are present in 1.5 L of 0.24M Na2SO4? M = n of solute L of solution “Triangle Trick” 0.24M = n 1.5 L 1.5 × × 1.5 Divide n 1.5L 0.24M 0.36 mol = n Multiply

32. Ex 3 A saline solution contains 0.90g NaCl in exactly 100.0 mL of solution. What is the molarity of the solution? 1. CONVERT GRAMS TO MOLES!! 2. CONVERT mL to L. M = n of solute L of solution What are you solving for? Molarity (M) M = 0.016 mol 0.1000 L 0.90g NaCl 1 mol NaCl = 0.016 mol 58g NaCl 100.0 mL 1L = 0.1000 L 1000mL Molarity = 0.16 M K h D M d c m L

33. 0.650M = n 2.40 L Ex 4 How many grams of solute are in 2.40L of 0.650M HClO2? M = n L 2.40 L× × 2.40 L 1.56 mol = n 68g HClO2 1.56 mol HClO2 = 106.08g HClO2 1 mol HClO2

34. Title: Molarity ProblemsYou do not have to write the problem. You MUST show your work. BOX in your answer! 1. A solution has a volume of 2.0L and contains 36.0g of glucose. If the molar mass of glucose is 180g, what is the molarity of the solution? 2. A solution has a volume of 250 mL and contains 0.70 mol NaCl. What is its molarity? 3. How many moles of ammonium nitrate are in 335 mL of 0.425M NH4NO3? 4. How many moles of solute are in 250 mL of 2.0M CaCl2? How many grams of CaCl2 is this?(molar mass = 110g)

35. VII. Concentrations of Solutions (cont.) B. Making Dilutions • You can make a solution less concentrated by diluting it with a solvent. • A dilution reduces the moles of solute per unit volume, however, the total moles of solute in solution does not change M1× V1 = M2 × V2 • Volumes can be in L or mL, as long as the same units are used for both V1 & V2

36. Ex 1 You need 150 mL of 0.40M NaCl and you have a 1.0M of NaCl solution. Calculate the volume of the NaCl solution. V1 = 150 mL of 0.40 NaCl M1 = 0.400M NaCl M2 = 1.0M NaCl Unknown = V2 M1× V1 = M2 × V2 0.400M × 150 mL = V2 1.0M × 60. mL = V2

37. Ex 2 What volume of 5.00M sulfuric acid is required to prepare 25.00L of 0.400M sulfuric acid? M1 = 5.00M V2 = 25.00L M2 = 0.400M M1× V1 = M2 × V2 5.0M × V1 = 0.40M × 25 L V1 = 2.0 L

38. Ex 3 How many milliliters of a stock solution of 2.00M MgSO4 would you need to prepare 100.0 mL of 0.400M MgSO4? M1 = 2.00M MgSO4 M2 = 0.400M MgSO4 V2 = 100 mL of 0.400M MgSO4 Unknown = V1 M1× V1 = M2 × V2 V1 = 0.400M ×100 mL 2.00M × V1= 20 mL

39. LEFT SIDETitle: Dilution ProblemsYou do not have to write the problem. You MUST show your work. BOX in your answer. 5. In making a dilution how many mL of 4.0M HCl are needed to make 250 mL of 0.30M HCl? 6. In making a dilution how many mL of a 3.00M NaBr solution are needed to make 175 mL of 0.400M NaBr? 7. How many milliliters of a stock solution of 2.00M MgSO4 would you need to prepare 100mL of 0.400M MgSO4? 8. How many milliliters of a stock solution of 4.00M KI would you need to prepare 250.0mL of 0.760M KI?

40. Adding to the mole road map!! Solutions ? mol 1 L 1 L ? mol Molarity: 5.0 M = 5.0 mol 1 L Molecule Formula unit Atoms (ions)

41. VIII. Using Molarity in Stoichiometry Zn + 2 HCl → ZnCl2 + H2 How many milliliters of 3.00M HCl are required to react with 17.35 g of zinc? 17.35 g Zn 1 mol Zn 65 g Zn 1 L HCl 3.00 mol HCl 1000 mL HCl 1 L HCl 2 mol HCl 1 mol Zn = 177.95 mL HCl

42. Using Molarity in Stoichiometry Cu + 2 AgNO3 → 2 Ag + Cu(NO3)2 How many grams of copper are required to react with 40.0 mL of 9.0M AgNO3? 9.0 mol AgNO3 1 L AgNO3 1 L AgNO3 1000 mL AgNO3 1 mol Cu 2 mol AgNO3 40.0 mL AgNO3 64 g Cu 1 mol Cu = 11.52 g Cu

43. IX. Percent Solutions If both solute and solvent are liquids the concentration of the solute is expressed as a percent. Percent by volume (%(v/v)) = volume of solute X 100 solution volume For solutions of solids dissolved in liquids is percent (mass/volume). Percent (mass/volume) (%(m/v)) = mass of solute (g) X 100 solution volume (mL)

44. IX. Percent Solutions • Ex 1 What is the percent by volume of ethanol (C2H6O), or ethyl alcohol, in the final solution when 85 mL of ethanol is diluted to a volume of 250 mL with water? % (v/v) = 85 mL ethanol X 100 250 mL solution = 34% ethanol (v/v)

45. IX. Percent Solutions • Ex 2 A solution of glucose (C6H12O6) contains 2.8 g in 20.0 mL of solution. What is the percent (m/v) of the solution? % (m/v) = 2.8g X 100 20.0 mL = 14% glucose (m/v)

46. Why do we get to make ice cream? X. Freezing-Point Depression • the difference in temperature between the freezing point of a solution and that of the pure solvent. • The salt depresses the freezing point of water because it disrupts the crystal formation of the water

47. Why do we get to make ice cream? When you add salt to the ice, it lowers the freezing point of the ice, so even more energy has to be absorbed from the environment in order for the ice to melt. This makes the ice colder than it was before, which is how your ice cream freezes Compounds that break into ions upon dissolving, like NaCl breaks into Na+ and Cl-, are better at lowering the freezing point than substances that don't separate into particles because the added particles disrupt the ability of the water to form crystalline ice. The more particles there are, the greater the disruption and the greater the impact on particle-dependent properties (colligative properties) like freezing point depression.

48. Ice Cream Lab 1. Put all ingredients on my desk. 2. While I am mixing the ingredients. You need to prepare your coffee can and freezer bag. 3. To prepare your coffee can, put a 5cm layer of ice then sprinkle 1cm of rock salt, then add another layer of ice then rock salt. (3/4th full, leave room to put in the ice cream bag) 4. Then you will get your freezer zip lock bag and duct tape the 3 side edges. 5. I will then put ice cream in your bag and you will seal the opening with duct tape. 6. Place the taped bag of ice cream inside the middle of the coffee can and duct tape the lid onto the coffee can. 7. Roll your coffee can outside for 20 minutes! 8. When complete, take bag out of can and rinse off bag. 9. Then enjoy ice cream!