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Chapter 13 Solutions Types of Solutions • A solution is a homogenous mixture of 2 or more substances • The solute is(are) the substance(s) present in the smaller amount(s) • The solvent is the substance present in the larger amount A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature. An unsaturated solution contains less solute than the solvent has the capacity to dissolve at a specific temperature. A supersaturated solution contains more solute than is present in a saturated solution at a specific temperature. Crystallization is the process in which dissolved solute comes out of solution and forms crystals Sodium acetate crystals rapidly form when a seed crystal is added to a supersaturated solution of sodium acetate. A Molecular View of the Solution Process Three types of interactions in the solution process: • solvent-solvent interaction • solute-solute interaction • solvent-solute interaction DHsoln = DH1 + DH2 + DH3 A Molecular View of the Solution Process DHsoln = DH1 + DH2 + DH3 • If the solute-solvent attraction is > the solute-solute and solvent-solvent attractions, then the solution process is favorable, or EXOTHERMIC (DHsoln < 0). • If the solute-solvent interaction is weaker than the solvent-solvent and solute-solute interactions, then the solution process is ENDOTHERMIC (DHsoln > 0). • Solution process is governed by 2 main factors: 1. Energy – determines if a solution process is endothermic or exothermic 2. Inherent tendency toward disorder in all natural events (entropy) – increase in randomness or disorder when solute and solvent molecules mix to form a solution * It is the increase in disorder of the system that favors the solubility of any substance, even if the solution process is endothermic Solubility – “Like Dissolves Like” Solubility is a measure of how much solute will dissolve in a solvent at a specific temperature. Two substances with similar intermolecular forces are likely to be soluble in each other. • non-polar molecules are soluble in non-polar solvents CCl4 in C6H6 • polar molecules are soluble in polar solvents C2H5OH in H2O • ionic compounds are more soluble in polar solvents NaCl in H2O or NH3 (l) • Two liquids are said to be Miscible if they are completely soluble in each other in all proportions. • Solvation is the process in which an ion or a molecule is surrounded by solvent molecules arranged in a specific manner. Is molecular iodine more soluble in water or in carbon disulfide? Carbon disulfide – a nonpolar solvent “like dissolves like”: nonpolar iodine will be soluble in nonpolar carbon disulfide Concentration Units The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. Percent by Mass Independent of temperature mass of solute % by mass = x 100% mass of solute + mass of solvent mass of solute = x 100% mass of solution Mole Fraction (X) moles of A X = A sum of moles of all components Concentration Units Molarity (M) moles of solute M = liters of solution Molality (m) Independent of temperature moles of solute m = mass of solvent (kg) What is the molality of a 5.86 M ethanol (C2H5OH) solution whose density is 0.927 g/mL? moles of solute moles of solute m = M = mass of solvent (kg) liters of solution Assume 1 L of solution: 5.86 moles ethanol = 270 g ethanol 927 g of solution (1000 mL x 0.927 g/mL) mass of solvent = mass of solution – mass of solute = 927 g – 270 g = 657 g = 0.657 kg moles of solute 5.86 moles C2H5OH m = = = 8.92 m mass of solvent (kg) 0.657 kg solvent The Effect of Temperature on Solubility Solid solubility and temperature solubility increases with increasing temperature solubility decreases with increasing temperature Fractional crystallization is the separation of a mixture of substances into pure components on the basis of their differing solubilities. Suppose you have 90 g KNO3 contaminated with 10 g NaCl. Fractional crystallization: 1. Dissolve sample in 100 mL of water at 600C 2. Cool solution to 00C 3. All NaCl will stay in solution (s = 34.2g/100g) 4. 78 g of PURE KNO3 will precipitate (s = 12 g/100g). 90 g – 12 g = 78 g Temperature and Solubility O2 gas solubility and temperature solubility usually decreases with increasing temperature The Effect of Pressure on the Solubility of Gases The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution (Henry’s law). c is the concentration (M) of the dissolved gas c = kP P is the pressure of the gas over the solution k is a constant for each gas (mol/L•atm) that depends only on temperature low P high P low c high c Pressure & Solubility of Gases Each gas has a different k value at a given temperature The amount of gas that will dissolve in a solvent depends on how frequently the gas molecules collide with the liquid surface and become trapped by the condensed phase Practical application: effervescence of a soft drink when the bottle cap is removed Calculate the molar concentration of oxygen in water at 25C for a partial pressure of 0.22 atm. The Henry’s law constant for oxygen is 1.3 x 10-3 mol/Latm. c=kP c = (1.3 x 10-3 mol/Latm)(0.22 atm) c = 2.9 x 10-4 mol/L How is oxygen highly soluble in blood when only sparingly soluble in water? Colligative Properties of Nonelectrolyte Solutions • Colligative properties are properties that depend only on the number of solute particles in solution and not on the nature of the solute particles. 1. Vapor-pressure lowering 2. Boiling-point elevation 3. Freezing-point depression 4. Osmotic pressure • Dealing with Dilute solutions (< or = 0.2 M) Colligative Properties 1. Vapor-Pressure Lowering • If a solute is nonvolatile (does not have a measurable vapor pressure), then the vapor pressure of its solution is always less than that of the pure solvent • Concentration of solute in solution determines relationship between vapor pressure of solution and vapor pressure of solvent 0 P 0 = vapor pressure of pure solvent P1 = X1 P 1 1 Raoult’s law X1 = mole fraction of the solvent If the solution contains only one solute: X1 = 1 – X2 0 0 P 1 - P1 = DP = X2 P 1 X2 = mole fraction of the solute Vapor-Pressure Lowering Vapor pressure of a solution is lower than that of pure solvent…Why? – Increasing disorder in a solution – Solvent molecules have less of a tendency to leave a solution than to leave the pure solvent to become vapor Volatile = has measurable vapor pressure – If both components of a solution are volatile, then the vapor pressure of the solution is the sum of the individual partial pressures Ideal Solution = any solution that obeys Raoult’s Law – Characteristics: Heat of solution, DHsoln = 0; volumes are additive An Ideal Solution 0 PA = XA P A 0 PB = XB P B PT = PA + PB 0 0 PT = XA P A + XB P B PT is greater than Intermolecular forces predicted by Raoults’s law between A & B molecules are weaker than those between A molecules & B molecules Greater tendency for these molecules to leave the solution Vapor pressure of solution > sum of vapor pressures predicted by Raoult’s Law Force Force Force Positive deviation < & A-B A-A B-B Heat of solution is positive (endothermic) A molecules attract B PT is less than molecules more predicted by Raoults’s law strongly than they do their own Vapor pressure of solution is < sum of vapor pressures predicted by Raoult’s Law Negative deviation Heat of solution is Force Force Force negative > & (exothermic) A-B A-A B-B Fractional Distillation Solution vapor pressure has direct bearing on Apparatus fractional distillation (procedure for separating liquid components of a solution based on different boiling points) At each step, the composition of the vapor in the column will be richer in the more volatile (lower bp) component The vapor that rises to the top of the column condenses & is collected in a receiving flask Colligative Properties 2. Boiling-Point Elevation DT = T – T 0 • Boiling-point elevation = boiling b b b point of the solution – boiling point of 0 T is the boiling point of pure solvent b the pure solvent •Solute must be nonvolatile T b is the boiling point of the solution 0 Tb > T b DTb > 0 DTb = Kb m m is the molality of the solution Kb is the molal boiling-point elevation constant (0C/m) for a given solvent (Table 12.2) Colligative Properties 3. Freezing-Point Depression 0 DTf = T f – Tf 0 • Freezing-point depression = freezing T is the freezing point of point of the pure solvent – freezing point f of the solution the pure solvent T f is the freezing point of the solution 0 T f > Tf DTf > 0 DTf = Kf m m is the molality of the solution Kf is the molal freezing-point depression constant (0C/m) for a given solvent (Table 12.2) Calculate the freezing point and boiling point of a solution containing 478 g of ethylene glycol in 3202 g water. DTf = Kf m 1 mol EG 478 g EG 7.70 mol EG 62.07g EG moles of solute 7.70 mol EG m 2.41 m kg solvent 3.202 kg water o o DTf K f m (1.86 C / m)(2.41m) 4.48 C Because pure water freezes at 0o C, the solution will freeze at - 4.48o C DTb = Kb m o o DTb Kbm (0.52 C / m)(2.41m) 1.25 C The solution will boil at (100 1.25)o C or 101.25o C.
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