The Solution Process. Solution - homogeneous mixture of solute and solvent. In solutions, intermolecular forces become rearranged.

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1 The Solution Process Solution - homogeneous mixture of solute and solvent. In solutions, intermolecular forces become rearranged.

2 Examples of solutions gas in gas e.g. air gas in liquid -- e.g. soda gas in solid -- e.g. gas on solid, catalyst liquid in liquid liquid in solid -- e.g. mercury amalgam solid in liquid solid in solid -- e.g. 14-karat gold, brass

3 Consider NaCl (solute) dissolving in water (solvent): Interruption of water H-bonds, NaCl Na + + Cl -, ion-dipole forces form: Na + -OH 2 and Cl - +H 2 O. If water is the solvent, we say the ions are hydrated.

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6 Energy Changes and Solution Formation 3 energy steps in forming a solution: separation of solute molecules ( H1), separation of solvent molecules ( H2), and formation of solute-solvent interactions ( H3).

7 H soln = H 1 + H 2 + H 3. H soln can be +ve or -ve depending on the intermolecular forces.

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10 Rule : polar solvents dissolve.? Non-polar solvents dissolve...?

11 Exercise: Why doesn t gasoline dissolve NaCl? Exercise: Why doesn t water and octane mix well (immiscible)? Remember: the resultant solution s interactions must be stronger than the interactions in the original substance

12 Solution Formation, Spontaneity, and Disorder When the energy of the system decreases (e.g. dropping a book and allowing it to fall to a lower potential energy), the process is spontaneous.

13 Example: a mixture of CCl 4 and C 6 H 14 is less ordered than the two separate liquids. Therefore, they spontaneously mix

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15 There are solutions that form by physical processes and those by chemical processes.

16 Consider: Ni(s) + 2HCl(aq) NiCl 2 (aq) + H 2 (g). 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 Consider: 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.

18 Saturated Solutions and Solubility Dissolve: solute + solvent Crystallization: solution solution. solute + solvent. Saturation: crystallization and dissolution are in equilibrium.

19 Solubility: amount of solute required to form a saturated solution. Supersaturation: reached when more solute is dissolved than in a saturated solution.

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21 Factors Affecting Solubility Solute-Solvent Interaction Miscible liquids: mix in any proportions. Immiscible liquids: do not mix. Intermolecular forces are important The more C atoms, the less the solubility in water.

22 The -OH groups in a molecule increase solubility in water. like dissolves like

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24 Which of these two would be more soluble in water?

25 Network solids do not dissolve. Why?

26 Pressure Effects Solubility of a gas in a liquid is a function of the pressure of the gas.

27 Pressure Effects

28 The higher the pressure, the more molecules of gas are close to the solvent

29 Henry s Law gives: S g kp g where: S g - solubility of a gas, k is a constant, and P g is the partial pressure of a gas

30 Example 27g of acetylene, C 2 H 2, dissolves in 1L of acetone at 1.0 atm pressure. If the partial pressure of acetylene is increased to 12 atm, what is the solubility in acetone? Solution: S 1 = kp 1 (1) S 2 = kp 2 (2) Ans: 3.2 x 10 2 g

31 Carbonated beverages are bottled with a partial pressure of CO 2 > 1 atm. What happens when a bottle is opened?

32 Temperature Effects As temperature increases, solubility of solids generally increases, e.g. sugar in warm water Sometimes, solubility decreases as temperature increases (e.g. Ce 2 (SO 4 ) 3 ).

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34 Gases - less soluble at high temperature Temperature Effects Thermal pollution in dams and rivers loss of O 2

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36 Ways of Expressing Concentration Definitions: Mass Percentage, ppm, and ppb mass % of component mass of component in solution total mass of solution 100

37 Example: How would you prepare 425 g of an aqueous solution containing 2.40% by mass of sodium acetate, NaC 2 H 3 O 3? Ans: Mass of NaC 2 H 3 O 3 = 10.2 g Mass of H 2 O = mass of solution - mass of NaC 2 H 3 O 3 = 415 g

38 Exercise: Concentrated aqueous nitric acid has 69.0% by mass of HNO 3 and has a density of 1.41 gcm -3. What volume of this solution contains 14.2 g of HNO 3? Ans: 14.6 cm 3

39 ppm of component massof component in solution total massof solution 6 10 Also mgl -1

40 ppb of component mass of component in solution total mass of solution 9 10 Also μgl -1

41 Exercise: Seawater contains g of dissolved oxygen, O 2, per litre. The density of seawater is 1.03 gcm -3. What is the concentration of oxygen, in ppm? Ans: 6.2 ppm

42 Mole Fraction, Molarity, and Molality Mole fraction of component moles of component in solution total moles of solution Molarity moles liters of solute solution

43 Molality, m moles solute kg of solvent Converting between molarity (M) and molality (m) requires density. Exercise: 0.2 mol of ethylene glycol is dissolved in 2000 g of water. Calculate the molality

44 Example: What is the molality of a solution containing 5.67 g of glucose, C 6 H 12 O 6 (M r = g), dissolved in 25.2 g of water? (Calc. the mole fractions of the components as well). Solution: Think about the solute!...glucose (express in moles) Think about the solvent!...water (express in kilograms) Ans: 1.25 m

45 Example: Converting molarity to molality An aqueous solution is M Pb(NO 3 ) 2. What is the molality of lead nitrate, Pb(NO 3 ) 2, in this solution? The density of the solution is g/ml. (Molar mass of Pb(NO 3 ) 2 = g) Solution: Mass of solution = density x volume Calculate mass of Pb(NO 3 ) 2, ie, moles x Mr Mass of H 2 O = mass of solution mass of Pb(NO 3 ) 2 Molality = m Pb(NO 3 ) 2

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47 Colligative Properties Colligative properties - depend only on the number of particles in solution and not on their identity. So NaCl(s) a + (aq) + Cl - (aq) K 2 SO 4 (s) C 12 H 22 O 11 (s) 2K + (aq) + SO 4 2- (aq) C 12 H 22 O 11 (aq)

48 Examining the effect of adding a non-volatile solute to a solvent on: 1. vapor pressure 2. boiling point 3. freezing point 4. osmosis

49 Examples are: anti-freeze in the radiator water in a car prevents freezing in winter and boiling in summer; snow is melted by adding salt on sidewalks and streets

50 Lowering Vapor Pressure VP lowering depends on the amount of solute.

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52 Raoult s law: Lowering Vapor Pressure Recall Dalton s Law: P soln = X solvent P o solvent P total = P A + P B + P C +.P N

53 Ideal solution - obeys Raoult s law Raoult s law is to solutions what the ideal gas law is to gases Raoult s law breaks down when the solvent-solvent and solutesolute intermolecular forces are greater than solute-solvent intermolecular forces For liquid-liquid solutions where both components are volatile, a modified form of Raoult s law applies: P total = P A + P B = X A P o A + X B P o B

54 Example: Predict the vapour pressure of a solution prepared by mixing 35 g solid Na 2 SO 4 (Mr = 142 g/mol) with 175 g water at 25 o C. The vapour pressure of pure water at 25 o C is torr. Ans: 22.1 torr

55 Exercise: The hydrocarbon limonene is the major constituent of lemon oil. A solution of limonene in 78.0 g of benzene had a vapour pressure of 90.6 mm Hg at 25 o C, and the vapour pressure of pure benzene at 25 o C is 95.2 mm Hg. What is its mass and molecular formula? Ans: C 10 H 16

56 As with gases, ideal behaviour for solutions is never perfectly achieved Nearly ideal behaviour is observed if solute-solute, solvent-solvent and solutesolvent interactions are very similar

57 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.

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59 Molal boiling-point-elevation constant, K b, expresses how much T b changes with molality, m: T b K b m

60 Freezing Point Depression T f K f m

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62 Freezing Point Depression

63 Example: How many grams of ethanol, C 2 H 5 OH, must be added to 37.8 g of water to give a freezing point of o C? Solution: Water is the solvent and ethanol the solute From table 13.4, Tf = 0.15 o C; K f for water is 1.86 o C/m T f K f m

64 Colligative properties of ionic solutions T f = ik f m where i is the no. of ions resulting from each formula unit

65 Example: Estimate the freezing point of a m aqueous solution of aluminium sulphate, Al 2 (SO 4 ) 3. Assume the value of i based on the formula of the compound. Ans: o C

66 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 through a semi-permeable membrane.

67 Eventually the pressure difference between the arms stops osmosis. Osmosis

68 Osmotic pressure,, is the pressure required to stop osmosis: V nrt n RT V MRT Isotonic solutions are solutions.? Osmosis

69 Hypotonic solutions are solutions.? Hypertonic solutions are solutions? Osmosis is spontaneous. Red blood cells are surrounded by semipermeable membranes.

70 Example: The formula for low-molecular weight starch is (C 6 H 10 O 5 ) n, where n averages 2x10 2. When g of starch is dissolved in 100 ml of water solution, what is the osmotic pressure at 25 o C? = atm

71 Exercise: Fish blood has an osmotic pressure equal to that of seawater. If seawater freezes at -2.3 o C, what is the osmotic pressure of the blood at 25 o C? Ans: 30 atm

72 Crenation: red blood cells placed in hypertonic solution (relative to intracellular solution); The cell shrivels or swells up?

73 Hemolysis: there is a higher solute concentration in the cell; What happens to the cell?

74 Osmosis

75 Hypertonic solution Hypotonic solution

76 To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic.

77 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).

78 Active transport is the movement of nutrients and waste material through a biological system. Active transport is not spontaneous.

79 Colloids A colloid is a dispersion of particles of one substance (the dispersed phase) throughout another substance or solution (the continuous phase) Examples.?? Butter, paint, milk, marshmallows, whipped cream, smoke, fog

80 The particle sizes range from ~1 x 10 3 pm to 2 x 10 5 pm in size

81 Although a colloid appears to be homogeneous because the dispersed particles are quite small, it can be distinguished from a true solution by its ability to scatter light This is called the effect?

82 Left: vessel containing colloid; Right: true solution

83 Aerosols liquid droplets or solid particles dispersed in a gas e.g. fog and smoke Emulsion liquid droplets dispersed throughout another liquid e.g. butterfat in milk Sol solid particles dispersed in a liquid e.g. AgCl(s) in H 2 O

84 Colloids in which the continuous phase in water can be hydrophilic (e.g. protein molecules) or hydrophobic colloids (Au particles in water).

85 How does soap stabilise oil in water? And how do we digest fats in our digestive systems?

86 Removal of colloidal particles Because of their small size, colloidal particles tend to be difficult to remove by processes such as filtration. Thus enlargement of particles by coagulation is required Heating or adding an electrolyte can cause coagulation

87 Heating increases collisions and hence particle size Electrolytes neutralise surface charges and reduce repulsions e.g. Alum in water purification, clay deposits in deltas Semi-permeable membranes can also be used to separate ions from colloidal particles (dialysis e.g. waste removal from blood by kidneys)