Bushra Javed Valencia College CHM 1046 Chapter 12 - Solutions

Transcription

1 Bushra Javed Valencia College CHM 1046 Chapter 12 - Solutions 1

2 Chapter 12 :Solutions Tentative Outline 1. Introduction to solutions. 2. Types of Solutions 3. Solubility and the Solution Process: Saturated, Unsaturated and Supersaturated Solutions 4. Solubility of Ionic Compounds 5. Effect of Temperature on the Solubility of Ionic Salts 6. Solubility of Molecular Solids & Liquids 7. Units of Concentration 8. Physical Behavior of Solutions: Colligative Properties: Vapor pressure, Boiling Point Elevation & Freezing Point Depression 2

3 What is a solution Solution A homogeneous mixture of two or more substances consisting of ions or molecules. A solution has two components, Solvent and Solute. Solvent determines the physical state of the solution & is present in greater quantity. Solute is the component in smaller amount 3

4 Type of Solution State of Matter Example Gas in gas Gas Air (O 2, N, Ar, CO 2 other gases) Gas in liquid liquid carbonated water (soda) CO 2 in water Gas in solid liquid Air in whipping cream Liquid in liquid liquid Gasoline (mixture of hydrocarbons) Liquid in solid solid Dental amalgam (mercury in silver) Solid in liquid liquid Sea waters (NaC1 and other salts in water Solid in solid solid 14-karat gold (gold in silver) Brass (Zn in copper) Stainless Steel (chromium in iron) 4

5 Solubility and the Solution Process: Saturated & Unsaturated Solutions: Saturated Solution: A saturated solution is one in which no more of the solute will dissolve at a given temperature. Unsaturated Solution: An unsaturated solution is one in which more of the solute could dissolve at a given temperature 5

6 Saturated & Unsaturated Solutions 6

7 Supersaturated Solution A supersaturated solution is a solution that contains more dissolved substance than a saturated solution does. This occurs when a solution is prepared at a higher temperature and is then slowly cooled. This is a very unstable situation, so any disturbance causes precipitation. 12 7

8 Saturated & Unsaturated Solutions Example 1: A small amount of a solid is added to water and, once equilibrium is reached, all the solid has dissolved. Which of the following statements is most likely to be true? a. The solution is supersaturated with solute. b. The solution is either unsaturated or supersaturated with solute. c. The solution is either saturated or supersaturated with solute. d. The solution is unsaturated with solute. 8

9 Solubility & The solution Process Solubility The amount of a substance that dissolves in a given quantity of water at a given temp to give a saturated solution. Solubility of NaCl = 36.0g/100ml at 20º and 1 atm Solubility of CO 2 = 0.161g/1.00ml at 20º and 1 atm 9

10 Solubility of Ionic Compounds Requires breaking of Lattice structure Lattice is a three-dimensional structure containing cations and anions held through strong forces of attraction Some ionic solids dissolve in water because water is a highly polar liquid. Strong ion-dipole interactions exist between water molecules and ions in aqueous solution. 10

11 Solubility of ionic solids Solubility of ionic solids depends on the following: 1. Hydration Energy: The attraction of ions for water molecules Hydration of ions favors dissolving. 2. Lattice Energy: Holds the ions together in the crystal lattice Works against the solution process. 11

12 Solubility of ionic solids 12

13 The force of attraction between water and both a cation and an anion is illustrated to the left with lithium fluoride, LiF.

14 Example 2 Solubility of ionic solids Which of the following concerning solubility and the solution process is/are correct? 1. Both hydration energy & lattice energy depend on the magnitude of ion charges and the size of the ions 2. An initially nonhomogeneous mixture of two miscible liquids, given enough time, will eventually form a solution as a result of random molecular motions 3. The dissolution of ionic compounds in water depends only on the hydration energy of the ions A. 1&2 only B. 1,2&3 C. 1& 3 only D. 3 only 14

15 Solubility of ionic solids Example 3 Which of the following pure liquids is the best solvent for sodium fluoride? a. CCl 4 (l) b. HCl(l) c. PCl 5 (l) d. CS 2 15

16 Effect of Temperature on solubility of ionic compounds The dissolving of ionic compounds in water is either exothermic or endothermic. If endothermic, solubility increases with an increase in the temperature. If exothermic, solubility decreases with an increase in temperature. Find the basis for instant hot and cold 16

17 In general, solubility depends on temperature.

18 Exothermic & Endothermic Solution Process 18

19 Exothermic & Endothermic Solution Process Example 4 Which of the following sets of conditions favors maximum solubility of an ionic solute in water? a. The magnitude of the lattice energy should be small, and the enthalpy of hydration of the ions should be large. b. The magnitude of the lattice energy should be small, and the enthalpy of hydration of the ions should be small as well. c. The magnitude of the lattice energy should be large, and the enthalpy of hydration of the ions should be small. 19

20 Solubility/Miscibility of Liquids A Liquid may dissolve in other liquid to produce a solution For molecular solutions, this can be summarized as Like dissolves like. Solutes dissolve in solvents that have the same type of intermolecular forces. 20

21 Solubility/Miscibility of Liquids An immiscible solute has different intermolecular forces from those of the solvent 21

22 Miscibility of Liquids through Hydrogen Bonds Water and methanol are miscible because the molecules in both the pure liquids and their mixtures form hydrogen bonds. 22

23 Example 6 Miscibility of Liquids In general, which of the following types of compounds would be the most soluble in carbon disulfide, CS 2? a. ionic b. polar molecular c. nonpolar molecular d. network covalent 23

24 Miscibility of Liquids Example 7 Why is gasoline immiscible in water 1. The attractions between gasoline molecules and water molecules are different from the gasolinegasoline molecular attractions and water-water molecules attractions. 2. London dispersion forces between gasoline molecules, and hydrogen bonds between water molecules are very similar. 3. London dispersion forces in gasoline are much stronger than the hydrogen bonds in water. 24

25 Miscibility of Liquids Example 8 Acetone is a versatile solvent. It has both a polar bond and a nonpolar hydrocarbon chain. Water is a polar compound. Would acetone dissolve in water? & Why Structural Formula of acetone. 25

26 Miscibility of Liquids Structural Formula of cyclohexane Example 9: Cyclohexane is a nonpolar organic compound. Would cyclohexane dissolve in acetone? 26

27 Solubility of Molecular Solids 27

28 Factors affecting the solubility of Molecular compounds Example 10 Which of the following sets of conditions favors maximum solubility of solute in solvent? a. The intermolecular forces between solute and solvent molecules are much stronger than the intermolecular forces between solute molecules or the intermolecular forces between solvent molecules. b. The intermolecular forces between solute and solvent molecules are much stronger than the intermolecular forces between solvent molecules. c. The intermolecular forces between solute and solvent molecules are much weaker than the intermolecular forces between solute molecules. 28

29 Solubility of molecular compounds Example 11: Which of the following pure liquids is the best solvent for carbon disulfide? a. C 6 H 6 (l) b. NH 3 (l) c. CH 3 OH(l) d. H 2 O(l) 29

30 Solution Strength:Concentration Units Dilute and concentrated are nonspecific terms. To know the exact amount of solute dissolved in a given amount of solvent, several concentration units are designed. 30

31 Solution Strength:Concentration Units The concentration of solutes can be quantitatively expressed in several ways: Molarity molality mass percentage of solute mol fraction 31

32 Molarity(M) Defined as moles of solute per liters of solution It is abbreviated as M: M = moles of solute liters of solution The most common way of expressing concentration in the chemistry laboratory. 32

33 Molarity Example 12: What is the molarity of NaI solution that contains 6.00g of NaI in 20.0mL of solution? a. 1.00M b M c. 2.00M d. 1.00M 33

34 molality Defined as number of moles of solute per kilogram of solvent Is abbreviated as m : m = moles of solute kilograms of solvent Has some advantage of over molarity Mass does not change with temperature while volume does. 34

35 molality Example 13 If 12.9 g of naphthalene, C 10 H 8, is dissolved in g of chloroform, CHCl 3, what is the molality of the solution? a m b m c m d m 35

36 Example 14: molality How many mol of urea (60.0g/mol)must be dissolved in 71.6g of water to give a 2.4m solution? a. 5.0x10^ 2 mol b. 0.17mol c. 1.4 x 10^ 2 mol d mol 36

37 Mass percentage of solute Mass percentage of solute is the percentage by mass of solute in a solution. Mass percentage of solute = grams of solute grams of solution 100% 37

38 Mass percentage of solute Example 15 What mass of a 33.0% by mass glucose, C 6 H 12 O 6, solution contains 54.0 g of glucose? a) 17.8 g b) 54.0 g c) 59.4 g d) g 38

39 Converting % mass into molarity Example 16 What is the molarity of a 40.0% by mass hydrochloric acid solution? The density of the solution is g/ml. a) 18.3 M b) 13.1 M c) M d) M 39

40 Converting mass percentage into Molarity Example 17 A sulfuric acid solution that is 65.0% H 2 SO 4 by mass has a density of 1.55g/mL at 20C. What is the molarity of sulfuric acid in solution? a. 3.5M b. 6.9M c. 10.3M d. 15.7M 40

41 Mol Fraction Mole fraction is the moles of component over the total moles of solution. It is abbreviated Χ. Χ = moles of solute total moles of solution 41

42 Mol Fraction Example 18 What is the mol fraction of water in a solution that contains 6.8mol of ethanol and 3.0 mol of water? a b c d

43 Mol Fraction Example 19 What is the mole fraction of urea in a solution that contains 3.4 mol of urea and 5.4 mol of water? a) 0.61 b) 0.39 c) 0.49 d)

44 The Colligative Properties Colligative properties of solutions are properties that depend on the concentration of the solute molecules or ions in solution but not on the chemical identity of the solute. 1.Vapor-pressure lowering 2.Boiling-point elevation 3.Freezing-point lowering 44

45 Vapor Pressure Vapor pressure is the pressure exerted by the gas in equilibrium with a solid or liquid in a closed container at a given temperature. A volatile substance has high vapor pressure while nonvolatile substance has negligible vapor pressure 45

46 Vapor Pressure of a solution 46

47 Raoult s Law Vapor pressure of a solution, P soln containing a nonvolatile solute equals the vapor pressure of pure solvent, P solv times the mole fraction of solvent, X solv : P soln = P solv X solv Thus lowering in vapor pressure is a colligative property.. 47

48 Vapor Pressure of a solution Example 20 When 1 mol of a nonvolatile nonelectrolyte is dissolved in 3 mol of a solvent, the vapor pressure of the solution compared with that of the pure solvent is a) 1/5. b) 1/4. c) 1/2. d) 3/4. 48

49 Phase diagram of a solution containing a nonvolatile solute 49

50 Elevation in Boiling point The boiling point of a solution is higher than the boiling point of pure solvent in the presence of a nonvolatile solute. The boiling-point elevation, T b, is given by the following equation: T b = mk b 50

51 Elevation in Boiling point The boiling point elevation of a solution relative to that of a pure solvent depends on the number of dissolved particles 51

52 Elevation in Boiling point Example 21 A solute added to a solvent raises the boiling point of the solution because a) the temperature to cause boiling must be great enough to boil not only the solvent but also the solute. b) the solute particles lower the solvent's vapor pressure, thus requiring a higher temperature to cause boiling. c) the solute particles raise the solvent's vapor pressure, thus requiring a higher temperature to cause boiling. d) the solute increases the volume of the solution, which requires an increase in the temperature 52

53 Elevation in Boiling point Example 22 The fact that the boiling point of a pure solvent is lower than the boiling point of a solution of the same solvent is a direct consequence of the a) vapor pressure of the solution being higher than the vapor pressure of the pure solvent. b) vapor pressure of the solution being lower than the vapor pressure of the pure solvent. c) osmotic pressure of the solvent being higher than the osmotic pressure of the solution. d) osmotic pressure of the solvent being lower than the osmotic pressure of the solution 53

54 Elevation in Boiling point Example 23 What is the boiling point change for a solution containing 0.347mol of naphthalene(a non volatile,non ionizing compound)in 250.g of liquid benzene?(kb =2.53C/m for benzene) a C b C c. 7.29C d C 54

55 Freezing Point Depression The freezing-point depression of a solution is a colligative property. The solvent molecules collide with crystals of solid solvent less frequently in the presence of a solute than they do in the pure solvent 55

56 Van t Hoff factor Van t Hoff factor i, is the number of particles a formula breaks up into ethylene glycol (C 2 H 6 O 2 ) (l) C 2 H 6 O 2(aq) (1 mole of particles). van t Hoff factor of C 2 H 6 O 2 =1 NaCl (s) Na + (aq) + Cl - (aq) (2 moles of particles) van t Hoff factor of NaCl = 2 CaCl 2(s) Ca 2+ (aq) + 2Cl - (aq) (3 moles of particles) van t Hoff factor of CaCl 2 =3 56

57 Van t Hoff factor Lowering in the freezing point by different solutes can be calculated through the following formula using van t Hoff factor: T = ik f m where T = Change in temperature in C i = van 't Hoff factor K f = molal freezing point depression constant in C kg/mol m = molality of the solute in mol solute/kg solvent. 57

58 Van t Hoff factor Example 24: For which of the following aqueous solutions would one expect to have the largest van t Hoff factor (i)? a m C 6 H 12 O 6 (glucose) b m NaCl c m NaCl d m K 2 SO 4 58

59 Example 25: Van t Hoff factor Which of the following solutes, dissolved in 1.0 kg of water, creates a solution that boils at the highest temperature? a mol HCl b mol HF c mol HClO 4 d mol H 2 SO 4 59

60 Freezing Point Depression Example 26: Which of the following will cause the calculated molar mass of a compound determined by the freezing-pointdepression method to be greater than the true molar mass? a. Water gets into the solvent after the freezing point of the pure solvent is determined. b. Some of the solute molecules break apart. c. The mass of solvent is smaller than that determined from the weighing. d. When the solute was added, some was spilled on the lab bench. 60

61 Freezing Point of a solution Example 27: What is the freezing point of an aqueous 1.18 m CaCl 2 solution? (K f for water is C/m.) Hint: Find T first. a. 2.2 C b. 2.2 C c. 6.6 C d. 6.6 C 61

62 Molar Mass from Freezing Point Example 28: Depression The molar mass of a substance may be determined by the freezing-point-depression technique. The minimum data required for the determination are a. T and K f. b. T, K f, and mass of solute. c. T, K f, mass of solute, and mass of solvent. d. T, K f, mass of solute, mass of solvent, and identity of solvent. 62

63 Molar Mass from Freezing Point Depression Example 29: If a 15.4-g sample of a nonelectrolyte is dissolved in g of water, the resulting solution will freeze at 0.93 C. What is the molar mass of the nonelectrolyte? (K f for water is C/m.) a. 150 g/mol b. 180 g/mol c g/mol d. 140 g/mol 63