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CHEMISTRY The Central Science 9th Edition Chapter 13 Properties of Solutions 11 Text, P. 417, review (Chapter 11) 21 13.1: The Solution Process • Solutions • homogeneous mixtures • Solution formation is affected by • strength and type of intermolecular forces • forces are between and among the solute and solvent particles 31 Text, P. 486 41 Hydration of solute • Attractive forces between solute & solvent particles are comparable in magnitude with those between the solute or solvent particles themselves • Note attraction of charges •What has to happen to: • Water’s H-bonds? • NaCl? •What intermolecular force is at work in solvation? Text, P. 486 Energy Changes and Solution Formation There are three energy steps in forming a solution: • the enthalpy change in the solution process is Hsoln = H1 + H2 + H3 • Hsoln can either be + or depending on the intermolecular forces 71 Text, P. 487 Text, P. 488 MgSO4 Hot Pack NH4NO3 Cold Pack • Breaking attractive intermolecular forces is always endothermic • Forming attractive intermolecular forces is always exothermic • To determine whether Hsoln is positive or negative, consider the strengths of all solute-solute and solutesolvent interactions: • H1 and H2 are both positive • H3 is always negative 91 • Rule: Polar solvents dissolve polar solutes Non-polar solvents dissolve non-polar solutes (like dissolves like) WHY? – If Hsoln is too endothermic a solution will not form – NaCl in gasoline: weak ion-dipole forces (gasoline is non-polar) – The ion-dipole forces do not compensate for the separation of ions 101 Solution Formation, Spontaneity, and Disorder • A spontaneous process occurs without outside intervention • When energy of the system decreases, 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 • Entropy 111 Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids •Therefore, they spontaneously mix even though Hsoln is very close to zero Text, P. 489 121 Solution Formation and Chemical Reactions • Example: Ni(s) + 2HCl(aq) NiCl2(aq) + H2(g) • When all the water is removed from the NiCl2 solution, no Ni is found only NiCl2·6H2O (a chemical reaction that results in the formation of a solution) • Water molecules fit into the crystal lattice in places not specifically occupied by a cation or an anion • Hydrates • Water of hydration • Think about it: What happens when NaCl is dissolved in 131 water and then heated to dryness? NaCl(s) + H2O (l) Na+(aq) + Cl-(aq) • When the water is removed from the solution, NaCl is found • NaCl dissolution is a physical process 141 • Sample problem # 3 151 13.2: 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 161 13.3: Factors Affecting Solubility 1. Solute-Solvent Interaction • “Like dissolves like” • Miscible liquids: mix in any proportions • Immiscible liquids: do not mix 171 Generalizations: • Intermolecular forces are important: • Water and ethanol are miscible • broken hydrogen bonds in both pure liquids are re-established in the mixture • The number of carbon atoms in a chain affects solubility: the more C atoms in the chain, the less soluble the substance is in water 181 Generalizations, continued: • The number of -OH groups within a molecule increases solubility in water • The more polar bonds in the molecule, the better it dissolves in a polar solvent (like dissolves like) • Network solids do not dissolve • the strong IMFs in the solid are not re-established in any solution 191 Text, P. 493 201 Read “Chemistry & Life”, P. 494 Fat soluble vitamin Water soluble vitamin 211 2. Pressure Effects • Solubility of a gas in a liquid is a function of the pressure of the gas 221 • High pressure means • More molecules of gas are close to the solvent • Greater solution/gas interactions • Greater solubility • If Sg is the solubility of a gas k is a constant Pg is the partial pressure of a gas then Henry’s Law gives: S g kPg Carbonated Beverages! 231 3. Temperature Effects Text, P. 497 • As temperature increases • Solubility of solids generally increases • Solubility of gases decreases • Thermal pollution 241 251 Figure 13.17, P. 497 • Sample problem # 17 261 13.4: Ways of Expressing Concentration • All methods involve quantifying amount of solute per amount of solvent (or solution) • Amounts or measures are masses, moles or liters • Qualitatively solutions are dilute or concentrated 271 • Definitions: mass of component in solution 100 1. mass % of component total mass of solution mass of component in solution ppm of component 10 6 total mass of solution mass of component in solution ppb of component 109 total mass of solution 281 2. moles of component in solution Mole fraction of component total moles of solution 3. moles solute Molarity liters of solution • Recall mass can be converted to moles using the molar mass 291 4. moles solute Molality, m kg of solvent • Converting between molarity (M) and molality (m) requires density • Molality doesn’t vary with temperature • Mass is constant • Molarity changes with temperature • Expansion/contraction of solution changes volume 301 Text, P. 501 • Sample Problems #31, 33, 37, 39, 41 321 13.5: Colligative Properties Colligative properties depend on quantity of solute particles, not their identity • Electrolytes vs. nonelectrolytes 0.15m NaCl 0.15m in Na+ & 0.15m in Cl- 0.30m in particles 0.050m CaCl2 0.050m in Ca+2 & 0.1m in Cl- 0.15m in particles 0.10m HCl 0.10m in H+ & 0.10m in Cl- 0.20m in particles 0.050m HC2H3O2 between 0.050m & 0.10m in particles 0.10m C12H22O11 0.10m in particles • Compare physical properties of the solution with those of the pure solvent 331 1. Lowering Vapor Pressure • Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid • Vapor pressure is lowered • Raoult’s Law: PA is the vapor pressure with solute PA is the vapor pressure without solute A is the mole fraction of solvent in solution A PA AP A 341 Ideal solution: one that obeys Raoult’s law • Raoult’s law breaks down (Real solutions) • Real solutions approximate ideal behavior when • solute concentration is low • solute and solvent have similar IMFs • Assume ideal solutions for problem solving 2. Boiling-Point Elevation • The triple point - critical point curve is lowered 351 • At 1 atm (normal BP of pure liquid) there is a lower vapor pressure of the solution • A higher temperature is required to reach a vapor pressure of 1 atm for the solution (Tb) • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m: Tb Kb m 361 Text, P. 505 3. Freezing Point Depression • The solution freezes at a lower temperature (Tf) than the pure solvent – lower vapor pressure for the solution • Decrease in FP (Tf) is directly proportional to molality (Kf is the molal freezing-point-depression constant): T f K f m 381 Text, P. 505 Applications: Antifreeze! 391 • Examples: # 45, 47, 49, 51 & 53 • A neat link 411 4. 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 • “Where ions go, water will flow” ~ Mrs. Moss 421 • Eventually the pressure difference between the arms stops osmosis Text, P. 507 431 • Osmotic pressure, , is the pressure required to stop osmosis: V nRT n RT V MRT • It is colligative because it depends on the concentration of the solute in the solvent 441 • Isotonic solutions: two solutions with the same separated by a semipermeable membrane • Hypertonic solution: a solution that is more concentrated than a comparable solution • Hypotonic solution: a solution of lower than a hypertonic solution • Osmosis is spontaneous • Read text, P. 508 – 509 for practical examples 451 • Examples: #57, 59 & 61 461 • There are differences between expected and observed changes due to colligative properties of strong electrolytes – Electrostatic attractions between ions – “ion pair” formation temporarily reduces the number of particles in solution – van’t Hoff factor (i): measure of the extent of ion dissociation 471 • Ratio of the actual value of a colligative property to the calculated value (assuming it to be a nonelectrolyte) – Ideal value for a salt is the # of ions per formula unit Tf(measured) i Tf(nonelectrolyte) Factors that affect i: •Dilution •Magnitude of charge on ions • lower charges, less 481 deviation • Sample Problem, # 63, 82 491 11.6: Colloids • Read Text, Section 13.6, P. 511 – 515 – Terms/Processes: • Tyndall effect • Hydrophilic • Hydrophobic • Adsorption • Coagulation 501 11.6: Colloids • Read Text, Section 13.6, P. 511 – 515 • Suspensions in which the suspended particles are larger than molecules • too small to drop out of the suspension due to gravity • Tyndall effect: ability of a colloid to scatter light • The beam of light can be seen through the colloid 511 Text, P. 512 521 Hydrophilic and Hydrophobic Colloids • “Water loving” colloids: hydrophilic • “Water hating” colloids: hydrophobic • Molecules arrange themselves so that hydrophobic portions are oriented towards each other 531 • Adsorption: when something sticks to a surface we say that it is adsorbed • Ions stick to a colloid (colloids appears hydrophilic) • Oil drop and soap (sodium stearate) • Sodium stearate has a long hydrophobic tail (Carbons) and a small hydrophilic head (-CO2-Na+) 541 Text, P. 514 Removal of Colloidal Particles • Coagulation (enlarged) until they can be removed by filtration • Methods of coagulation: – heating (colloid particles are attracted to each other when they collide) – adding an electrolyte (neutralize the surface charges on the colloid particles) 561 End of Chapter 13 Properties of Solutions 571