Sheets Page 1 Lecture 30 Lecture 30: Solutions 4 Read: BLB 13.1–13.3 HW: BLB 13:7,15,21,23,33 Sup13:1–7 Know: • how solutions form • solubility & saturation • factors affecting solubility Need help?? Get help!! TAs in CRC (211 Whitmore) and SI—hours on Chem 110 website; my office hours (Mon 12:30-2 & Tues 10:30-12 in 324 Chem Bldg [or 326 Chem]) Bonus deadline for BST #9: Solutions & dilutions, April 2 Review chemical nomenclature (e.g., BST #5, Lecture 8 …); & memorize those strong acids & bases (BLB Table 4.2) Check out the grade-u-lator @ http://courses.chem.psu.edu/chem110/spring/grade.htm Exam #3: Monday, April 6 @ 6:30 pm; Sign up for the conflict exam, if needed. Last day to sign up is Wed, April 1. No text-programmable calculators (PDAs, iPods, etc). Bring PSU ID and several pencils; Late drop deadline: Friday, April 10 @ 11:59 pm via elionSheets Page 2 Lecture 30 The solution process • solvation: solutions form when the attractive IM forces between solute & solvent are comparable in magnitude (& nature) to those that exist between solute–solute or solvent–solvent alone • hydration: solvation in water • for solvation to occur: 1. break solute–solute interactions ⇒ 2. break solvent–solvent interactions ⇒ 3. solute–solvent interactions form ⇒Sheets Page 3 Lecture 30 !H1ΔH2 ΔH3Sheets Page 4 Lecture 30 The solution process (cont.) • processes occur spontaneously when: energy is released (exothermic), ΔH –; & disorder increases (entropy), ΔS ↑ • these processes are linked: a process can be endothermic when the increase in entropy (disorder) is large enough; this is true for any type of process (dissolution, reaction, etc.) • forming a solution always increases disorder because of mixing, thus spontaneity depends on ΔH • ΔHsoln = ΔH1 + ΔH2 + ΔH3 • ΔHsoln: overall enthalpy change in forming a solution • a solution will form in most cases unless solute–solute or solvent–solvent interactions are too strong, relative to the solute–solvent interactionsSheets Page 5 Lecture 30 Dynamic equilibrium • dynamic equilibrium: no change, BUT change is occurring on a molecular level; this equilibrium can be a physical change (e.g., vapor pressure or solvation) or a chemical change (more in Chap 15) forward rate = backward rate NOTE: vs. solute+solventsolutiondissolvecrystallizeSheets Page 6 Lecture 30 Solubility • solubility: amount of a substance that can be dissolved in solvent under given conditions (e.g., T) • saturated solution: solution that contains the concentration of solute under the given conditions; a solution in dynamic equilibrium with undissolved solute; additional solute will NOT dissolve if added to the solution • unsaturated solution: solution that contains than maximum concentration of solute; more solute can dissolve if added to the solution • supersaturated solution: solution that contains than maximum concentration of solute; more than the equilibrium amount; solution is unstable (that is, NOT at equilibrium)Sheets Page 7 Lecture 30 Solubility generalization • in general: like dissolves like • polar solvents dissolve & solutes • nonpolar solvents dissolve solutes • consider IM forces that are broken & formed; if they are the ~same, then dissolution is likely solubility in water alcohol at 25°C (g/100g of H2O) CH3OH ∞ CH3CH2OH ∞ CH3CH2CH2OH ∞ CH3CH2CH2CH2OH 8.06 CH3CH2CH2CH2CH2OH 2.82 CH3CH2CH2CH2CH2CH2OH 0.62 • as length of hydrocarbon chain ↑, solubility of alcohol in water ; why is this the case???Sheets Page 8 Lecture 30 To (phase) separate or not to (phase) separate…, that is the question a mol solute/100 g H2O at 20°C • miscible: 2 substances dissolve in all proportions • immiscible: 2 substances that do NOT dissolve in each other to a significant extent • examples NaCl in water? KCl in hexane? sugar (polar) in water? Ni metal in water? NH3 in water or hexane? in hexane; in waterSheets Page 9 Lecture 30 Example: Which one of the following will be most soluble in benzene [C6H6 (l)]? A. H2O(l) B. CH3OH(l) C. HCl(l) D. CH3CH2OH(l) E. heptane (C7H16) (l)Sheets Page 10 Lecture 30 Factors that affect solubility • other factors: T, P • temperature can either increase or decrease solubility, depending on ΔHsoln for ionic solids, solubility for gases, solubility • Henryʼs law: relationship between solubility of gases ↑ with its partial pressure; solids & liquids not affected; why pop (umm, I mean soda) fizzes Cg = KH Pg • Cg [Cg on data sheet; Sg in BLB!!] = solubility of gas in solution phase • KH = Henryʼs law constant; different for each solute/solvent pair; varies with temperature • Pg = partial pressure of gas (the solute) over solutionSheets Page 11 Lecture 30 Le Chatelierʼs principle • a dynamic equilibrium tends to oppose any change in conditions (more about this in Chap 15!) so how do T & P changes affect solubility, anyway? 1. pressure: as P increases, the system tries to reduce P, so more gas dissolves 2. temperature: as T increases, the system tries to reduce T, using the solution process to take up heat • if ΔHsoln < 0 (heat released, exothermic), e.g., solute comes out of solution (gases) • if ΔHsoln > 0 (heat taken up, endothermic), solute dissolves more (ionic solids)* *ΔHsoln > 0 at saturation for most ionic solidsSheets Page 12 Lecture 30 More examples To increase the solubility of O2 in water, need to do which of the following A. increase T B. decrease T C. increase PO2 D. decrease PO2 How will the solubility of KClO3 be affected by an increase in pressure?Sheets Page 13 Lecture 30 Before next class: Read: BLB 13.5–13.6 HW: BLB 13:9,58,61,67,69,75 Sup13:12–18 Know: • colligative properties vapor pressure lowering boiling point elevation freezing point depression osmotic pressure • colloids Answers: p. 9: E p. 12: top, B,C; bottom, KClO3 solubility is not affected by
View Full Document
Unlocking...