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Part One: Introduction A. Terminology. 1. Solution CHAPTER 12: SOLUTIONS Part One: Introduction A. Terminology. 1. Solution - homogeneous mixture of substances in which no settling occurs. 2. Consists of a solvent and one or more solutes. 3. Solute = is a substance being dissolved, and usually present in small amounts. 4. Solvent = medium in which the solutes are dissolved. Solvent is usually the most abundant species in the mixture. 5. Solubility = the amount of a given solute that dissolves in a given quantity of solvent. NaCl in water, e.g., solubility of NaCl in water = 3.6 g/mL. 6. Two fluids that mix in all proportions to form a solution are said to be miscible. (alcohol and water). Oil and water would be said to be immiscible. Figure 12.1 B. Importance. 1. Extremely important in all life – e.g., body fluids are solutions. 2. A solid, liquid, or gas can act as either solvent or solute. Usually the solvent is a liquid. 3. Sea water is an aqueous solution of many salts & some gases. 4. Tap water is an aqueous solution with trace amounts of various ions, molecular compounds, and gases. (Distilled water is pure). Chapter 12 Page 1 5. Air is a solution of gases. 6. Solids solutions are called alloys. One example is steel, which is mostly Fe with small amounts of additives like carbon, chromium, titanium, etc. Another example is dental fillings are solid solutions called by the special name amalgams = metals dissolved in Hg. Part Two: The Solution Process A. Solubility and Saturated Solutions 1. Solution formation is a process of dynamic equilibrium. Example, dissolving table salt in water: NaCl(s) Na+(aq) + Cl-(aq) Figure 12.2 Chapter 12 Page 2 Ions are continually leaving the solid crystal into the solution and returning to the crystal, until eventually: a) all the solid is dissolved, OR b) the solution reaches its saturation point, at which some solid remains. Which case pertains depends on the solubility of NaCl in water. Figure 12.3 B. Factors that influence solubility. 1. General rule = “like dissolves like.” e.g., oil dissolves in gasoline because both are hydrocarbons, which are non-polar molecules water, however, is a polar molecule, such that oil does not dissolve in it, but alcohols do (which are also polar). Therefore, intermolecular forces are involved. 2. There is some natural tendency of all substances to mix, based on tendency toward a state of increased disorder (ENTROPY CHANGE) 3. This is balanced by the tendency of a system to have the lowest energy possible (ENERGY CHANGE). In other words, is heat absorbed or released in forming a solution? Release of energy (exothermic) favors solution formation. Chapter 12 Page 3 The Entropy effect and the Energy effect combine to determine the solubility. A process is favored by: a. decrease in the energy of the system, exothermic process. b. increase in the disorder, or randomness of the system. 4. Energy change is called the heat of solution, ΔHsolution. Depends on how strongly solute and solvent particles interact. If ΔHsolution is negative → heat liberated (exothermic) If ΔHsolution is positive → heat absorbed (endothermic) The Entropy effect always works in favor of forming the solution. 5. Relative strengths of these interactions determine the extent of solubility (refer to figure below) a. Strong solvent-solute attractions favor solubility. b. Weak solvent-solvent attractions favor solubility. c. Weak solute-solute attractions favor solubility. 6. Solution formation is always accompanied by an increase in the disorder of both solvent and solute, always favorable to solubility. C. Dissolving Solids in Liquids, viewed as a stepwise process (Section 12.2) Chapter 12 Page 4 1. Take an ionic solid. Must overcome crystal lattice energy (solute-solute interactions): Requires large input of energy (endothermic) unfavorable to solution. 2. Must overcome solvent-solvent interactions. If solvent is water, must break Hydrogen bonds: Requires input of energy (endothermic) unfavorable. 3. Must “solvate” the solute particles: Exothermic, favorable to solution, especially ionic compounds dissolving in H2O. Figure 12.8 Chapter 12 Page 5 4. Combined Steps 2 and 3 is called “solvation,” or in case of water, “hydration.” 5. Overall energy involved: ΔHsolution = Steps 2 & 3 Step 1 ΔHsolution = (heat of solvation) - (crystal lattice energy) OR (heat of hydration) favorable unfavorable Figure 12.10 6. Energy is not the only factor. Entropy increase upon solution is a favorable factor which can sometimes overcome unfavorable energy factor if latter is not too large. + - 7. Example: NaCl(s) + H2O(aq) → Na (aq) + Cl (aq) ΔHsolution is (+), endothermic (unfavorable). Still occurs due to slightly larger favorable entropy consideration. 8. Ionic compounds which have very large crystal lattice energies will be insoluble in H2O because solvation and entropy can’t overcome energy holding lattice together. BaSO4 insoluble! 9. Some ionic solids dissolve with release of heat (exothermic) - both energy and entropy favor solution process. CaCl2, Na2SO4, ... Chapter 12 Page 6 10. Nonpolar solids (e.g., paraffin) do not dissolve in H2O because interactions with water (hydration) is very weak compared to water-water interactions. Weak interaction Strong Hydrogen Weak interaction btwn. in nonpolar solid bonding solvent- solvent & nonpolar (should be easy to solvent interactions. solute. Nothing gained break up). (difficult to break up) here except entropy of HOWEVER... mixing. Result: Water “prefers” itself rather than a non-ionic or non-polar solute. Called the “hydrophobic effect.” D. Dissolving Liquids in Liquids – Molecular solutions (Miscibility). (Section 12.2) 1. Same factors involved. Here, though, a solid does not need to be broken up, so no lattice energy to overcome. 2. Case 1: Polar or Hydrogen-bonding solute with polar or Hydrogen bonding solvent. Methanol dissolves in water. H-bonding. Moderately H-bonding. Moderately H-bonding. Nothing strong solute-solute strong solvent-solvent lost, entropy of interactions. interactions. mixing favors. 3. Case 2: Nonpolar solute with polar or Hydrogen-bonding solvent. CCl4 does not dissolve in H2O. Weak solute-solute H-bonding. Moderately Weak solute-solvent interactions. (only strong solvent-solvent. interactions. (loss of London forces) solvent-solvent int.) (hydrophobic effect) Chapter 12 Page 7 4. Case 3: Nonpolar solute and nonpolar solvent. CCl4 does dissolve in benzene. Weak solute-solute Weak solvent-solvent Weak solute-solvent interactions. (only interactions. (only (London). Entropy London forces) London forces) of mixing favors this. 5. Generalization - “like dissolves like.” E. Gases in Liquids. (Section 14.4) 1. Same factors as before. 2. Polar gases dissolve in polar liquids, nonpolar gases dissolve in nonpolar liquids. 3. CO2 and O2 are nonpolar, so only slightly dissolve in H2O. 4. CO2 more so because of ionization: CO2(g) + H2O(l) H2CO3(aq) carbonic acid + - H2CO3 (aq) H (aq) + HCO3 (aq) 5. HCl and NH3 are polar covalent gases, dissolve strongly in H2O. Part Three: Effects of Environmental Variables on Solubility (Section 12.3) A. Temperature Effect. (Section 12.3) 1. Explained by LeChatelier’s Principle = when a stress is applied to a system in equilibrium, the system responds in a way that relieves the stress and restores a new equilibrium. 2. When heat of solution is exothermic: Solute + Solvent → Solution + Heat T↑ causes shift back to the left, solubility↓. Chapter 12 Page 8 3. When heat of solution is endothermic: Solute + Solvent + Heat → Solution T↑ causes shift to the right, solubility↑. 4. NaCl solubility↑ as T↑; Na2SO4 solubility↓ as T↑. Figure 12.12 5. O2 dissolves in H2O exothermically: T↑ O2 solubility↓ . Basis of “thermal pollution.” B. Pressure Effect. 1. Henry’s Law = Increase of partial pressure of a gas over a liquid increases amount of gas dissolved by a direct proportion 2. Example: Carbonated beverage stored with pressurized CO2. 3. Henry’s Law: S = kH P (solubility of = kH (partial pressure of gas in liquid) gas above liquid) kH = Henry’s Law constant Chapter 12 Page 9 Part Four: Expressions of Concentration (Section 12.4) A. Molarity. 1. Molarity = moles of solute Liters of solution B. Mass Percentage of solute. 1. Mass % solute = mass of solute x 100% Mass of solution 2. Example problem: A certain solution of an alcohol and water is 2.0 M alcohol and has a density of 0.95 g/mL. The alcohol has a molar mass of 126 g/mol. What is the mass percent alcohol in this solution? C. Molality. 1. m = moles of solute kg of solvent 2. Example: What is the molality of a solution in which 1.0 gram glucose (C6H12O6) has been added to 100 mL of H2O? 3. Note that: molality ≈ molarity for very dilute aqueous solutions. Why? Because under these conditions: (Liters of solution) ≈ (kg of solvent) D. Mole Fraction. 1. XA = moles of component A total moles of all components Chapter 12 Page 10 2. Unitless. 3. Example: What is Xglucose and Xwater in previous problem? Part Five: Colligative Properties of Solution (Sections 12.5 through 12.8) A. Definition. 1. Colligative property = a physical property of a solution that depends upon the number, but not the identity, of solute particles in a given amount of solvent. 2. Examples: a. vapor pressure lowering b. boiling point elevation c. freezing point depression d. osmotic pressure 3. Depends on: total molar concentration of solute particles, not on their identity. B. Vapor Pressure Lowering - Raoult’s Law. (Section 12.5) 1. A solution containing a nonvolatile solute has a lower vapor pressure than the pure solvent. (“nonvolatile solute” means solute has no vapor pressure of its own) 2. Physical Picture: 3. V.P. lowering depends on fraction of surface blocked by solute particles, and thus, mole fraction of solute and solvent, hence: Chapter 12 Page 11 4. Raoult’s Law: Psolvent = Xsolvent P°solvent ↑ ↑ ↑ v.p. of mole v.p.
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