Presentation is loading. Please wait.

Presentation is loading. Please wait.

Prentice Hall © 2003Chapter 13 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition David P. White.

Similar presentations


Presentation on theme: "Prentice Hall © 2003Chapter 13 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition David P. White."— Presentation transcript:

1 Prentice Hall © 2003Chapter 13 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition David P. White

2 Prentice Hall © 2003Chapter 13 A solution is a homogeneous mixture of solute (present in smallest amount) and solvent (present in largest amount). Solutes and solvent are components of the solution. In the process of making solutions with condensed phases, intermolecular forces become rearranged. The Solution Process

3 Prentice Hall © 2003Chapter 13 The Solution Process

4 Prentice Hall © 2003Chapter 13 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. The Solution Process

5

6

7 Prentice Hall © 2003Chapter 13 Energy Changes and Solution Formation 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. The Solution Process

8 Prentice Hall © 2003Chapter 13

9 Prentice Hall © 2003Chapter 13 Energy Changes and Solution Formation Breaking attractive intermolecular forces is always endothermic. Forming attractive intermolecular forces is always exothermic. The Solution Process

10 Prentice Hall © 2003Chapter 13 Energy Changes and Solution Formation 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 ). The Solution Process

11

12 Prentice Hall © 2003Chapter 13 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. The Solution Process

13 Prentice Hall © 2003Chapter 13 Energy Changes and Solution Formation –Water in octane: water has strong H-bonds. There are no attractive forces between water and octane to compensate for the H-bonds. The Solution Process

14 Prentice Hall © 2003Chapter 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). The Solution Process

15 Prentice Hall © 2003Chapter 13 Solution Formation, Spontaneity, and Disorder If the process leads to a greater state of disorder, then the process is spontaneous. 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. The Solution Process

16 Prentice Hall © 2003Chapter 13 Solution Formation, Spontaneity, and Disorder The Solution Process

17 Prentice Hall © 2003Chapter 13 Solution Formation and Chemical Reactions Example: a mixture of CCl 4 and C 6 H 14 is less ordered 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. The Solution Process

18 Prentice Hall © 2003Chapter 13 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. The Solution Process

19 Prentice Hall © 2003Chapter 13 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. Saturated Solutions and Solubility

20 Prentice Hall © 2003Chapter 13

21 Prentice Hall © 2003Chapter 13 Solute-Solvent Interaction 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. Factors Affecting Solubility

22 Prentice Hall © 2003Chapter 13 Solute-Solvent Interaction The number of -OH groups within a molecule increases solubility in water. 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. Factors Affecting Solubility

23 Prentice Hall © 2003Chapter 13 Solute-Solvent Interaction Factors Affecting Solubility

24 Prentice Hall © 2003Chapter 13 Solute-Solvent Interaction Factors Affecting Solubility

25 Prentice Hall © 2003Chapter 13 Solute-Solvent Interaction Network solids do not dissolve because the strong intermolecular forces in the solid are not re-established in any solution. Pressure Effects Solubility of a gas in a liquid is a function of the pressure of the gas. Factors Affecting Solubility

26 Prentice Hall © 2003Chapter 13 Pressure Effects Factors Affecting Solubility

27 Prentice Hall © 2003Chapter 13 Pressure Effects 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. If S g is the solubility of a gas, k is a constant, and P g is the partial pressure of a gas, then Henrys Law gives: Factors Affecting Solubility

28 Prentice Hall © 2003Chapter 13 Pressure Effects Carbonated beverages are bottled with a partial pressure of CO 2 > 1 atm. As the bottle is opened, the partial pressure of CO 2 decreases and the solubility of CO 2 decreases. Therefore, bubbles of CO 2 escape from solution. Factors Affecting Solubility

29 Prentice Hall © 2003Chapter 13 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 ). Factors Affecting Solubility

30

31 Prentice Hall © 2003Chapter 13 Temperature Effects Experience tells us that carbonated beverages go flat as they get warm. Therefore, gases get less soluble as temperature increases. Thermal pollution: if lakes get too warm, CO 2 and O 2 become less soluble and are not available for plants or animals. Factors Affecting Solubility

32

33 Prentice Hall © 2003Chapter 13 Mass Percentage, ppm, and ppb 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: Ways of Expressing Concentration

34 Prentice Hall © 2003Chapter 13 Mass Percentage, ppm, and ppb Parts per million (ppm) can be expressed as 1 mg of solute per kilogram of solution. –If the density of the solution is 1g/mL, then 1 ppm = 1 mg solute per liter of solution. Parts per billion (ppb) are 1 g of solute per kilogram of solution. Ways of Expressing Concentration

35 Prentice Hall © 2003Chapter 13 Mass Percentage, ppm, and ppb Mole Fraction, Molarity, and Molality Recall mass can be converted to moles using the molar mass. Ways of Expressing Concentration

36 Prentice Hall © 2003Chapter 13 Mole Fraction, Molarity, and Molality We define Converting between molarity (M) and molality (m) requires density. Ways of Expressing Concentration

37

38 Prentice Hall © 2003Chapter 13 Colligative properties depend on quantity of solute molecules. (E.g. freezing point depression and melting point elevation.) Lowering 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. Colligative Properties

39 Prentice Hall © 2003Chapter 13 Lowering Vapor Pressure Colligative Properties

40 Prentice Hall © 2003Chapter 13 Lowering Vapor Pressure Raoults Law: P A is the vapor pressure with solute, P A is the vapor pressure without solvent, and A is the mole fraction of A, then Recall Daltons Law: Colligative Properties

41 Prentice Hall © 2003Chapter 13 Lowering Vapor Pressure Ideal solution: one that obeys Raoults law. Raoults law breaks down when the solvent-solvent and solute-solute intermolecular forces are greater than solute-solvent intermolecular forces. Boiling-Point Elevation Goal: interpret the phase diagram for a solution. Non-volatile solute lowers the vapor pressure. Therefore the triple point - critical point curve is lowered. Colligative Properties

42

43 Prentice Hall © 2003Chapter 13 Boiling-Point Elevation At 1 atm (normal boiling point of pure liquid) there is a lower vapor pressure of the solution. Therefore, a higher temperature is required to teach a vapor pressure of 1 atm for the solution ( T b ). Molal boiling-point-elevation constant, K b, expresses how much T b changes with molality, m: Colligative Properties

44 Prentice Hall © 2003Chapter 13 Freezing Point Depression At 1 atm (normal boiling point of pure liquid) there is no depression by definition When a solution freezes, almost pure solvent is formed first. –Therefore, the sublimation curve for the pure solvent is the same as for the solution. –Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for the solution. Colligative Properties

45 Prentice Hall © 2003Chapter 13 Freezing Point Depression The melting-point (freezing-point) curve is a vertical line from the triple point. The solution freezes at a lower temperature ( T f ) than the pure solvent. Decrease in freezing point ( T f ) is directly proportional to molality (K f is the molal freezing-point-depression constant): Colligative Properties

46 Prentice Hall © 2003Chapter 13 Freezing Point Depression Colligative Properties

47

48 Prentice Hall © 2003Chapter 13 Osmosis Semipermeable membrane: permits passage of some components of a solution. Example: cell membranes and cellophane. 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. Colligative Properties

49 Prentice Hall © 2003Chapter 13 Osmosis Eventually the pressure difference between the arms stops osmosis. Colligative Properties

50 Prentice Hall © 2003Chapter 13 Osmosis Osmotic pressure,, is the pressure required to stop osmosis: Isotonic solutions: two solutions with the same separated by a semipermeable membrane. Colligative Properties

51 Prentice Hall © 2003Chapter 13 Osmosis Hypotonic solutions: a solution of lower than a hypertonic solution. Osmosis is spontaneous. Red blood cells are surrounded by semipermeable membranes. Colligative Properties

52 Prentice Hall © 2003Chapter 13 Osmosis Crenation: –red blood cells placed in hypertonic solution (relative to intracellular solution); –there is a lower solute concentration in the cell than the surrounding tissue; –osmosis occurs and water passes through the membrane out of the cell. –The cell shrivels up. Colligative Properties

53 Prentice Hall © 2003Chapter 13 Osmosis Colligative Properties

54 Prentice Hall © 2003Chapter 13 Osmosis Hemolysis: –red blood cells placed in a hypotonic solution; –there is a higher solute concentration in the cell; –osmosis occurs and water moves into the cell. –The cell bursts. To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic. Colligative Properties

55 Prentice Hall © 2003Chapter 13 Osmosis –Cucumber placed in NaCl solution loses water to shrivel up and become a pickle. –Limp carrot placed in water becomes firm because water enters via osmosis. –Salty food causes retention of water and swelling of tissues (edema). –Water moves into plants through osmosis. –Salt added to meat or sugar to fruit prevents bacterial infection (a bacterium placed on the salt will lose water through osmosis and die). Colligative Properties

56 Prentice Hall © 2003Chapter 13 Osmosis Active transport is the movement of nutrients and waste material through a biological system. Active transport is not spontaneous. Colligative Properties

57 Prentice Hall © 2003Chapter 13 Colloids are suspensions in which the suspended particles are larger than molecules but too small to drop out of the suspension due to gravity. Particle size: 10 to 2000 Å. There are several types of colloid: –aerosol (gas + liquid or solid, e.g. fog and smoke), –foam (liquid + gas, e.g. whipped cream), –emulsion (liquid + liquid, e.g. milk), –sol (liquid + solid, e.g. paint), Colloids

58 Prentice Hall © 2003Chapter 13 –solid foam (solid + gas, e.g. marshmallow), –solid emulsion (solid + liquid, e.g. butter), –solid sol (solid + solid, e.g. ruby glass). Tyndall effect: ability of a Colloid to scatter light. The beam of light can be seen through the colloid. Colloids

59 Prentice Hall © 2003Chapter 13 Colloids

60 Prentice Hall © 2003Chapter 13 Hydrophilic and Hydrophobic Colloids Focus on colloids in water. Water loving colloids: hydrophilic. Water hating colloids: hydrophobic. Molecules arrange themselves so that hydrophobic portions are oriented towards each other. If a large hydrophobic macromolecule (giant molecule) needs to exist in water (e.g. in a cell), hydrophobic molecules embed themselves into the macromolecule leaving the hydrophilic ends to interact with water. Colloids

61

62 Prentice Hall © 2003Chapter 13 Hydrophilic and Hydrophobic Colloids Typical hydrophilic groups are polar (containing C-O, O- H, N-H bonds) or charged. Hydrophobic colloids need to be stabilized in water. Adsorption: when something sticks to a surface we say that it is adsorbed. If ions are adsorbed onto the surface of a colloid, the colloids appears hydrophilic and is stabilized in water. Consider a small drop of oil in water. Add to the water sodium stearate. Colloids

63 Prentice Hall © 2003Chapter 13 Hydrophilic and Hydrophobic Colloids Colloids

64 Prentice Hall © 2003Chapter 13 Hydrophilic and Hydrophobic Colloids Sodium stearate has a long hydrophobic tail (CH 3 (CH 2 ) 16 -) and a small hydrophobic head (-CO 2 - Na + ). The hydrophobic tail can be absorbed into the oil drop, leaving the hydrophilic head on the surface. The hydrophilic heads then interact with the water and the oil drop is stabilized in water. Colloids

65 Colloids

66 Prentice Hall © 2003Chapter 13 Hydrophilic and Hydrophobic Colloids Most dirt stains on people and clothing are oil-based. Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water. Bile excretes substances like sodium stereate that forms an emulsion with fats in our small intestine. Emulsifying agents help form an emulsion. Colloids

67 Prentice Hall © 2003Chapter 13 Removal of Colloidal Particles Colloid particles are too small to be separated by physical means (e.g. filtration). Colloid particles are coagulated (enlarged) until they can be removed by filtration. Methods of coagulation: –heating (colloid particles move and are attracted to each other when they collide); –adding an electrolyte (neutralize the surface charges on the colloid particles). Colloids

68 Prentice Hall © 2003Chapter 13 Removal of Colloidal Particles Dialysis: using a semipermeable membranes separate ions from colloidal particles. Colloids

69 Prentice Hall © 2003Chapter 13 End of Chapter 13 Properties of Solutions


Download ppt "Prentice Hall © 2003Chapter 13 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition David P. White."

Similar presentations


Ads by Google