1 CHAPTER 10 Solutions

2 Types of Solutions Solution - homogeneous mixture of 2 or more substances »solvent - dissolving medium »solute - dissolved species

3 Spontaneity of the Dissolution Process Assume solvent is a liquid Major factors that affect dissolution of solutes »change of energy content,  solution –exothermic favors dissolution –endothermic does not favor dissolution »change in disorder, or randomness,  S mixing –increase in disorder favors dissolution –increase in order does not favor dissolution Best conditions for dissolution »exothermic & disordered

4 Spontaneity of the Dissolution Process Disorder in mixing is very common »helps dissolution Factors that affect  solution solute-solute attractions requires absorption of E to overcome solvent-solvent attractions requires absorption of E to overcome solvent-solute attractions releases energy

5 Spontaneity of the Dissolution Process

6 The Solution Process A solution is a homogeneous mixture of solute (present in smallest amount) and solvent (present in largest amount). In the process of making solutions with condensed phases, intermolecular forces become rearranged. Consider NaCl (solute) dissolving in water (solvent): »the water H-bonds have to be interrupted, »NaCl dissociates into Na + and Cl -, »ion-dipole forces form: Na + …  -OH 2 and Cl - …  +H 2 O. »We say the ions are solvated by water. »If water is the solvent, we say the ions are hydrated.

7 The Solution Process

8 There are three energy steps in forming a solution: »separation of solute molecules (  H 1 ), »separation of solvent molecules (  H 2 ), and formation of solute-solvent interactions (  H 3 ). We define the enthalpy change in the solution process as  H soln =  H 1 +  H 2 +  H 3.  H soln can either be positive or negative depending on the intermolecular forces. Energy Changes and Solution Formation

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10 Energy Changes and Solution Formation Breaking attractive intermolecular forces is always endothermic. Forming attractive intermolecular forces is always exothermic. To determine whether  H soln is positive or negative, we consider the strengths of all solute-solute and solute-solvent interactions: »  H 1 and  H 2 are both positive. »  H 3 is always negative. »It is possible to have either  H 3 > (  H 1 +  H 2 ) or  H 3 < (  H 1 +  H 2 ).

11 Energy Changes and Solution Formation

12 Energy Changes and Solution Formation Examples: »NaOH added to water has  H soln = kJ/mol. »NH 4 NO 3 added to water has  H soln = kJ/mol. “Rule”: polar solvents dissolve polar solutes. Non- polar solvents dissolve non-polar solutes. Why? »If  H soln is too endothermic a solution will not form. »NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions. »Water in octane: water has strong H-bonds. There are no attractive forces between water and octane to compensate for the H-bonds.

13 Solution Formation, Spontaneity, and Disorder A spontaneous process occurs without outside intervention. When energy of the system decreases (e.g. dropping a book and allowing it to fall to a lower potential energy), the process is spontaneous. Some spontaneous processes do not involve the system moving to a lower energy state (e.g. an endothermic reaction). If the process leads to a greater state of disorder, then the process is spontaneous.

14 Solution Formation, Spontaneity, and Disorder Example: a mixture of CCl 4 and C 6 H 14 is less ordered than the two separate liquids. Therefore, they spontaneously mix even though  H soln is very close to zero. There are solutions that form by physical processes and those by chemical processes.

15 Solution Formation, Spontaneity, and Disorder

16 Solution Formation and Chemical Reactions Consider: Ni(s) + 2HCl(aq)  NiCl 2 (aq) + H 2 (g). Note the chemical form of the substance being dissolved has changed (Ni  NiCl 2 ). When all the water is removed from the solution, no Ni is found only NiCl 2.6H 2 O. Therefore, Ni dissolution in HCl is a chemical process.

17 Solution Formation and Chemical Reactions

18 Solution Formation and Chemical Reactions Example: NaCl(s) + H 2 O (l)  Na + (aq) + Cl - (aq). When the water is removed from the solution, NaCl is found. Therefore, NaCl dissolution is a physical process.

19 Dissolution of Solids in Liquids crystal lattice energy - energy absorbed when a mole of formula units of a solid is separated into its constituent ions (molecules or atoms for nonionic solids) in the gas phase »measure of attractive forces in solid »crystal lattice energy increases as charge density increases energy required to overcome London forces, dipole-dipole or H-bonding

20 Dissolution of Solids in Liquids dissolution is a competition between 1solute -solute attractions crystal lattice energy 2solvent-solvent attractions H-bonding, etc.

21 Dissolution of Solids in Liquids energy is released when solute particles are dissolved »energy of solvation »hydration energy (in water) look at dissolution of CaCl 2.

22 Dissolution of Solids in Liquids Ca OH 2 H2OH2O H2OH2O 2+ Cl - O H O H O H O

23 Dissolution of Solids in Liquids molar energy of hydration - energy absorbed when one mole of formula units becomes hydrated

24 Dissolution of Solids in Liquids hydration energy increases with increasing charge density

25 Dissolution of Liquids in Liquids (Miscibility) Most polar liquids are miscible with other polar liquids “like dissolves like” rule »methanol, CH 3 OH, very soluble in water

26 Dissolution of Liquids in Liquids (Miscibility) nonpolar liquids are miscible with other nonpolar liquids “like dissolves like” rule »nonpolar molecules “slide” in between each other

27 Rates of Dissolution and Saturation Finely divided solids dissolve more rapidly than large crystals granulated sugar vs. sugar cubes look at a single cube of NaCl enormous increase in surface area dissolves faster NaCl Breaks up many smaller crystals

28 Rates of Dissolution and Saturation saturated solutions have established an equilbrium between dissolved and undissolved solutes »air with 100% humidity »liquid solutions with solids supersaturated solutions have higher concentrations of dissolved solutes than saturated

29 Saturated Solutions and Solubility Dissolve: solute + solvent  solution. Crystallization: solution  solute + solvent. Saturation: crystallization and dissolution are in equilibrium. Solubility: amount of solute required to form a saturated solution. Supersaturated: a solution formed when more solute is dissolved than in a saturated solution.

30 Factors Affecting Solubility Solute-Solvent Interactions Polar liquids tend to dissolve in polar solvents. Miscible liquids: mix in any proportions. Immiscible liquids: do not mix. Intermolecular forces are important: water and ethanol are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture. The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water.

31 Factors Affecting Solubility

32 Factors Affecting Solubility The number of -OH groups within a molecule increases solubility in water.

33 Factors Affecting Solubility Generalization: “like dissolves like”. The more polar bonds in the molecule, the better it dissolves in a polar solvent. The less polar the molecule the less it dissolves in a polar solvent and the better is dissolves in a non-polar solvent. Network solids do not dissolve because the strong intermolecular forces in the solid are not re-established in any solution.

34 Effect of Pressure on Solubility Solubility of a gas in a liquid is a function of the pressure of the gas. The higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution. »Therefore, the higher the pressure, the greater the solubility. »The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility.

35 Pressure Effects

36 Pressure Effects Pressure changes have little or no effect on solubility of liquids and solids in liquids Pressure changes have large effects on the solubility of gases in liquids »why carbonated drinks fizz when opened »cause of several scuba diving related problems including the “bends”

37 Pressure Effects Phenomenon described by Henry’s Law

38 Effect of Temperature on Solubility ionic solids that dissolve endothermically »dissolution enhanced by heating ionic solids that dissolve exothermically »dissolution decreased by heating

39 Temperature Effects Experience tells us that sugar dissolves better in warm water than cold. As temperature increases, solubility of solids generally increases. Sometimes, solubility decreases as temperature increases (e.g. Ce 2 (SO 4 ) 3 ).

40 Temperature Effects

41 Temperature Effects

42 Temperature Effects Experience tells us that carbonated beverages go flat as they get warm. Gases are less soluble at higher temperatures. Thermal pollution: if lakes get too warm, CO 2 and O 2 become less soluble and are not available for plants or animals.

43 Ways of Expressing Concentration All methods involve quantifying amount of solute per amount of solvent (or solution). Generally amounts or measures are masses, moles or liters. Qualitatively solutions are dilute or concentrated. Definitions:

44 Ways of Expressing Concentration Mole Fraction, Molarity, and Molality Recall mass can be converted to moles using the molar mass. Recall

45 Ways of Expressing Concentration We define Converting between molarity (M) and molality (m) requires density. In dilute solutions, molarity and molality are nearly equal.

46 Ways of Expressing Concentration

47 Molality & Mole Fraction Calculate the molarity and the molality of an aqueous solution that is 10.0% glucose, C 6 H 12 O 6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C 6 H 12 O 6 = 180 g

48 Molality & Mole Fraction Calculate the molality and the molarity of an aqueous solution that is 10.0% glucose, C 6 H 12 O 6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C 6 H 12 O 6 = 180 g

49 Molality & Mole Fraction Calculate the molality and the molarity of an aqueous solution that is 10.0% glucose, C 6 H 12 O 6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C 6 H 12 O 6 = 180 g

50 Molality & Mole Fraction Calculate the molality and the molarity of an aqueous solution that is 10.0% glucose, C 6 H 12 O 6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C 6 H 12 O 6 = 180 g

51 Molality & Mole Fraction Calculate the molality of a solution that contains 7.25 g of benzoic acid C 6 H 5 COOH, in 200 mL of benzene, C 6 H 6. The density of benzene is g/mL. 1 mol C 6 H 5 COOH = 122 g

52 Molality & Mole Fraction Calculate the molality of a solution that contains 7.25 g of benzoic acid C 6 H 5 COOH, in 200 mL of benzene, C 6 H 6. The density of benzene is g/mL. 1 mol C 6 H 5 COOH = 122 g

53 Molality & Mole Fraction Mole fraction »number of moles of one component per moles of all the components of the solution »literally is a fraction using moles as the numerator and denominator »in a 2 component solution Mole fraction of component A - X A

54 Molality & Mole Fraction Mole fraction of component B - X B

55 Molality & Mole Fraction What are the mole fractions of glucose and water in a 10.0% glucose solution?

56 Molality & Mole Fraction What are the mole fractions of glucose and water in a 10.0% glucose solution?

57 Molality & Mole Fraction What are the mole fractions of glucose and water in a 10.0% glucose solution?

58 Molality & Mole Fraction now let’s calculate the mole fractions

59 Colligative Properties of Solutions Colligative properties »solution properties that depend solely on the number of particles dissoved in the solution and not the kinds of particles dissolved »physical property of solutions »four common types of colligative properties ¶ vapor pressure lowering · freezing point depression ¸ boiling point elevation ¹ osmotic pressure

60 Lowering the Vapor Pressure Non-volatile solvents reduce the ability of the surface solvent molecules to escape the liquid. Therefore, vapor pressure is lowered. The amount of vapor pressure lowering depends on the amount of solute.

61 Lowering the Vapor Pressure

62 Lowering of Vapor Pressure & Raoult’s Law Addition of a nonvolatile solute to a solution lowers the vapor pressure of the solution »simply due to fewer solvent molecules at surface »solute molecules occupy some of the spaces Raoult’s Law describes this effect in ideal solutions

63 Lowering of Vapor Pressure & Raoult’s Law Lowering of vapor pressure,  P solvent, is defined as:

64 Lowering of Vapor Pressure & Raoult’s Law Since X solvent + X solute = 1, we can derive

65 Lowering of Vapor Pressure & Raoult’s Law

66 Raoult’s Law Ideal solution: one that obeys Raoult’s law. Raoult’s law breaks down when the solvent- solvent and solute-solute intermolecular forces are greater than solute-solvent intermolecular forces.

67 Boiling Point Elevation Addition of a nonvolatile solute to a solution raises the boiling point of the solution above that of the pure solvent. »vapor pressure is lowered - Raoult’s Law »T must be raised to make vapor pressure equal to atmospheric pressure Amount that T is elevated is determined by the number of moles of solute dissolved in solution.

68 Boiling Point Elevation

69 Boiling Point Elevation Boiling point elevation relationship is:

70 Boiling Point Elevation What is the normal boiling point of a 2.50 m glucose, C 6 H 12 O 6, solution?

71 Boiling Point Elevation What is the normal boiling point of a 2.50 m glucose, C 6 H 12 O 6, solution?

72 Freezing Point Depression Addition of a nonvolatile solute to a solution lowers the freezing point of the solution relative to the pure solvent.

73 Freezing Point Depression Relationship for freezing point depression is:

74 Freezing Point Depression Notice the similarity of the two relationships for freezing point depression & boiling point elevation. Fundamentally, it is the same phenomenon. »differences are the size of the effect »reflected in the sizes of the constants, K f & K b Easily seen on a phase diagram for a solution.

75 Freezing Point Depression

76 Freezing Point Depression Calculate the freezing point of a 2.50 m aqueous glucose solution.

77 Freezing Point Depression Calculate the freezing point of a solution that contains 8.50 g of benzoic acid (C 6 H 5 COOH, MW = 122) in 75.0 g of benzene, C 6 H 6.

78 Freezing Point Depression

79 Determination of M.W. by Freezing Point Depression Freezing point depression depends on 2 things ¶size of K f for a given solvent –many of these values are well known ·molal concentration –# of moles of solute –kg of solvent If K f and kg of solvent are known, as is often the case in an experiment, then we can determine # of moles of solute and use it to determine the molecular weight.

80 Determination of M.W. by Freezing Point Depression A 37.0 g sample of a new covalent compound, a nonelectrolyte, was dissolved in 200 g of water. The resulting solution froze at C. What is the molecular weight of the compound?

81 Determination of M.W. by Freezing Point Depression A 37.0 g sample of a new covalent compound, a nonelectrolyte, was dissolved in 200 g of water. The resulting solution froze at C. What is the molecular weight of the compound?

82 Osmosis Semipermeable membrane: permits passage of some components of a solution. Osmosis: the movement of a solvent from low solute concentration to high solute concentration. There is movement in both directions across a semipermeable membrane. As solvent moves across the membrane, the fluid levels in the arms becomes uneven.

83 Osmosis

84 Osmosis Osmotic pressure, , is the pressure required to stop osmosis:

85 Osmosis Isotonic solutions: two solutions with the same  separated by a semipermeable membrane. Hypotonic solutions: a solution of lower  than a hypertonic solution. Osmosis is spontaneous. Red blood cells are surrounded by semipermeable membranes.

86 Group Question Medicines that are injected into humans, intravenous fluids and/or shots, must be at the same concentration as the existing chemical compounds in blood. For example, if the medicine contains potassium ions, they must be at the same concentration as the potassium ions in our blood. Such solutions are called isotonic. Why must medicines be formulated in this fashion?