Chem 206 Week 4 Lecture Notes (Kotz Ch.11-13) PDF

CHEM 206 – Dr.Rogers - Week 4 (of 12) ASSIGNED READINGS: NOW: finish Kotz 10th Ch.11 & start Ch.13 NEXT: continue Ch.13 SELF-STUDY: SLIDES WITH SS ON THEM MUST BE STUDIED IN DETAIL ON YOUR OWN. IN CLASS, THOSE WILL BE CLARIFIED, +/- EXERCISES...

CHEM 206 – Dr.Rogers - Week 4 (of 12) ASSIGNED READINGS: NOW: finish Kotz 10th Ch.11 & start Ch.13 NEXT: continue Ch.13 SELF-STUDY: SLIDES WITH SS ON THEM MUST BE STUDIED IN DETAIL ON YOUR OWN. IN CLASS, THOSE WILL BE CLARIFIED, +/- EXERCISES, BUT NOT “TAUGHT”. (1) 11.6 Properties of Liquids: bulk behaviour dictated by IMFs Total strength of intermolecular forces: 1. Types of forces: H-bonding, dipole-dipole, van der Waals 2. Extent of molecule over which they can act H2 H2 H2 H2 H2 H2 H2 H2 H2 Weak individual forces… H3C C C C C C C C C C but strong IMFs in total. C C C C C C C C C CH3 H2 H2 H2 H2 H2 H2 H2 H2 H2 (many points of contact between large molecules) grease (C20H42) actually a solid at RT Characteristic physical properties of LIQUIDS: 1. Ease of vaporization  vapour pressure, boiling point (in detail) 2. Surface tension, viscosity, capillary action (overview only)  Surface Tension = strength of liquid’s “skin”  Viscosity = resistance to flow molecules must slide wrt each other H2O(l)  Meniscus shape: cohesive vs adhesive forces  Capillary Action = liquid crawling up other Hg(l) (2) surfaces (tube, paper…)  Some molecules escape liquid: “vapour pressure”  as T  For liquid in a sealed container: Zumdahl Fig. 11.10 liquid & vapour eventually EQUILIBRATE See Kotz Fig.11.11 & see Appendix G  rate of evaporation = rate of condensation e.g., H2O(l) H2O(g) Practical note: Pvap ∝ conc. of vapour over liquid  use P = (n/V)RT  Pvap at given T = characteristic of liquid # of molecules → T1 Fewer  If liquid has high Pvap , what can we infer? T2 escape  liquid easily evaporated: “volatile” IMFs at  caused by weak IMFs (low ∆Hovap) lower T K Fig.  Clausius-Clapeyron: ln Pvap= -(∆H ovap /RT) + c 12.13 find Pvap at many Ts to determine ∆H ovap kinetic E → ∆Hovap (3) Relative humidity: % of water’s eqm Pvap in air (at given T) 100% humidity: vapour P of water in air = its eqm Pvap (= curve below) if “ > 100%” : it would rain (i.e., approach eqm from Q > K side) if < 100%: spilled water would evaporate (i.e., Q < K) Eqm P°H2O: H2O(l) H2O(g) ON OWN: Will all the water Kotz evaporate? (Kotz 9th Ch.11 p.432) Fig.11.12 Suppose 0.50 g of pure water is sealed in an evacuated 5.0 L flask and heated to 60°C. Will any liquid water be left in the flask? Appendix G: P°H O = 149.4 mm Hg 2 HINT: how much H2O(l) needs to evaporate to yield the eqm vapour P? NOTE: liquid will evaporate until its vapour generates the liquid’s eqm vapour P at that T, unless the liquid runs out before that P is reached. Boiling point: T at which vapour P = Pext When a liquid’s vapour P equals the atmospheric P, vapour pushes up as much as air pushes down on liquid  vapour bubbles form inside bulk liquid = BOILING KITCHEN WISDOM: leave lid on pot & water boils faster  vapours escape less  P→Pext in shorter time (same T). Kotz Fig.11.12 Kotz Fig.11.14 ON YOUR OWN: Why does water boil at a lower temperature at high elevations? Notice: diethyl ether has heaviest molecules, but is most volatile. (5) = Proof that IM forces type/strength (not MM!) determines Pvap. just for interest’s sake (not tested): critical T & P Strange behaviour at very high temperature & pressure SS  Critical Temperature & Pressure Kotz Fig.11.15 above critical point, distinction between liquid & vapour phases ceases to exist molecules pack close like in liquid, but with such high KE that move freely like in gas supercritical fluid = very-high-P gas with liquid-like density unexpected: powerful solvents CO2 used to extract caffeine (6) ON YOUR OWN: Predicting physical properties via IMFs Isomers with different connectivity of atoms normally have different physical properties (e.g., mp, bp, solubility), because of differences in their intermolecular forces. Which (pure) C2H6O isomer should have the higher bp, and why? Ethanol Dimethyl ether A. Dimethyl ether; d.-d. forces stronger than i.d.-d. forces B. Dimethyl ether; H-bonding stronger than d.-d. forces C. Ethanol; H-bonding in both, but stronger in ethanol D. Ethanol; H-bonding stronger than d.-d. forces Approach: Visualize which parts of each molecule interact in which way. List all active forces, then compare the “strongest” ones. (7)(Both molecules have similar # & types of atoms  similar van der Waals forces.)  CHAPTER 13 Solutions & their Behaviour Chapter Goals: (see text for full list) 13.1 Units of concentration  Understand the energetics of the 13.2 Solution process solution process & explain the basis of “like dissolves like” 13.3 Factors affecting solubility:  Calculate & use the various solution pressure & temperature concentration units 13.4 Colligative properties  Understand how changes in temperature & pressure affect solubility 13.5 Colloids (at level of notes)  Understand & use Henry’s law for gas solubility & Raoult’s law for vapour P RECOMMENDED PROBLEMS:  Understand & use the colligative Ch.13 #20, 38, 51, 57, 63, 70, properties of solutions 72, 86, 92, 94, 98, 100, 104, 106  Use colligative properties to determine the molar mass of a solute  Recognize the properties and importance of colloids If we want to get a feeling for how much solute should dissolve, we need to consider what’s going on… (8) STARTING WITH SECTIONS 13.2, 13.3… Energetics of the solution process: focusing on enthalpy… Z’s Fig.11.1 SOLUTION FORMATION: hypothetical steps 1) separate solute molecules (overcome intermol. forces) ∆H1 endo 2) separate solvent molecules (overcome intermol. forces) ∆H2 endo 3) combine them (results in new intermolecular forces) ∆H3 exo overall: “heat of solution” ∆Hsoln Could be Should ionic or consider covalent entropic changes too… but Ch.19 > Ch.14, so they don’t Determining factor: how favourable is hydration/solvation (step 3) ? ∆H3 often called enthalpy of hydration (or solvation) For ionic solutes, depends on: ion charge, ion size, solvent’s polarity (9) Factors affecting solubility: chemical structure H OH O H Polar solutes are “hydrophilic” H C H C C H C OH (= water-loving) HO O C C O HO C H …because water does not sacrifice IMF H H C OH C C strength when surrounds polar solute H C OH HO OH i.e., H2O↔H2O interactions similar to H C OH H2O↔POLAR-molecule interactions Vitamin C H (ascorbic acid) sugars Alcohols O δ - H δ+ δ- H O H H H δ+O H O C C δ H+ Hδ + H O H O H H C C H H H H H H H H Solvent SURROUNDS solute  touches EVERY surface of molecule H-bonding interactions: δ − O  δ + H, δ + H  δ − O (shown as ) d-d interactions: δ− O  δ+ C (not shown here) (10) van der Waals forces: δ − O  CHs, δ + H  CHs (not shown here) Non-polar molecules are “hydrophobic” (= water-fearing) H2O↔H2O interactions stronger than H2O↔NP-molecule interactions  Enthalpic sacrifice: to surround NP solute, water would need to give up strong IMFs with itself to make weaker IMFs with NP solutes  Water excludes most NP solutes: H2O prefers itself; NP substances float/sink Non-polar solvents DISSOLVE non-polar solutes CH3 CH3 CH3 H hydrocarbons H H H C C C C C C H H2 H2 H2C C C C C C OH C H H H H HC CH H3C C C H2C C C C CH 3 CH3 C HC CH H2 H2 H2 CH3 Vitamin A (retinol) C H polar OH group not enough to overwhelm non-polar majority carbon H2 H2 H2 H2 H2 H2 H2 H2 H2 tetrachloride H 3C C C C C C C C C C C C C C C C C C C CH 3 H2 H2 H2 H2 H2 H2 H2 H2 H2 grease (C20H42) (11)  Enthalpic thinking: “like dissolves like” & liquid miscibility  Consider interactions between solute & solvent molecules: if solvent-solvent interactions much stronger than solvent-solute int’ns, solubility will be LOW…and substances might not mix at all.  Liquid-liquid mixtures: if solution forms: liquids are “miscible” if separate layers result: “immiscible” Kotz 8th Fig.12.5 MISCIBLE: solute↔solvent interactions similar in IMMISCIBLE: strength to solute↔solute & solvent↔solvent int’ns Solvent excludes solute (12) Like dissolves like: applies to solid solutes too Kotz Fig.13.5 I2(s) = nonpolar Low solubility in water (polar) More soluble in nonpolar solvent BEFORE MIXING AFTER MIXING Must consider nature of the solid’s “building blocks”: ionic solids often quite soluble in water  hydrated ions polar molecular solids quite soluble in water  hydrated molecules nonpolar molecular solids can be soluble in nonpolar solvents  solvated molecules network solids: 3D array of covalent bonds (one GIANT molecule) e.g., diamond, silica (sand) (see Kotz p.457 briefly)  essentially not soluble in any solvent (13) CLICKER Q: Finding a good solvent In which solvent do you expect SO2(g) to be highly soluble? O O S A. Pentane - CH3CH2CH2CH2CH3 (l) B. Carbon tetrachloride - CCl4 (l) C. Carbon disulfide - CS2 (l) S C S D. Methanol – CH3OH (l) In which solvent would LiCl be most soluble? What type of interactions would be dominant? (14) Dissolving a SALT: ions become “hydrated” (solvated) Water molecules surround or “hydrate” the ions! anion cation Energetics of system: see Kotz Fig. 13.6 Enthalpic view: ionic bonds (∆Hlattice) → ion−dipole interactions (∆Hhydration) Entropic view: crystal + random H2Os → solvated ions + organized H2Os OVERALL: actually quite difficult to predict ∆Hsoln and ∆Ssoln here… (15) CLICKER Q: Explaining an observation Sof NaOH(s) = 64.5 J/K vs. Sof NaOH(aq) = 48.1 J/K Which of the statements below (some of which are entirely false!) provides a plausible explanation for the DECREASE in entropy upon dissolution of sodium hydroxide? In solution, A. the ions constrain water molecules from free motion. B. the ions are more tightly bound to each other when dissolved. C. some Na+ and OH- ions exist as clumps of ions. D. the Na+ ion is reduced to Na(s), which has a lower S°. (16) 13.3 Effect of T  Rate of dissolution (always) Temperature Solubility? …not always ! TRENDS FOR AQUEOUS SOL’NS: Dissolved gases: solubility with ↑ T Kotz dissolving gas entropically unfavourable ! Fig.13.11 thus: ∆Gdissolution becomes more +ve as T ↑ To de-gas a solution, warm it up. Dissolved liquids: hard to predict dissolving is entropically neutral (or close) Dissolved solids: solubility often with T To crystallize out solutes, cool down. Less soluble in hot water: EXCEPT: some ions (≥ 2±) bind H2O so CaCO3, CaSO4 & many other strongly that ∆Sdiss < 0 (sol’y at high T) Mg2+ & Ca2+ salts in hard water HARD WATER SCALE BUILDS UP AT HIGH T To predict T-dependence: use entropy analysis (∆G = ∆H – T∆S) the Ch.13 way (enthalpy (& Le Châtelier’s principle) isn’t actually the cause! Summary: Temperature dependence of solubility (13.3) ∆Sdissolution in water More soluble in… Solute (& description of why) hot or cold water ? gas usually < 0 (loss of disorder) cold water gas before, liquidy after liquid hard to predict (~ 0) can’t generalize liquid before, liquid after molecular usually > 0 (gain of disorder) hot water solid solid before, liquidy after (& H2Os loosely bound to solutes can still move as freely as before) ionic solid usually > 0 (gain of disorder) usually hot water solid before, liquidy after EXCEPT: highly charged ions can restrict movement of bound-H2Os some: cold water! (“hydration sphere”) so much that net ∆Sdiss < 0 ! (18) 13.5: When a solution is NOT really a solution… Solutions = homogeneous mixtures: CLEAR (= transparent) THUS, by definition: no chunks! NOT cloudy (= turbid) Colloids A suspension = suspensions A solution of particles… CLOUDINESS Fig.13.18 is due to: “the Tyndall effect” = light scattering off Laser  of suspended particles (even really small ones) 13.5 is filled with practical applications -but not enough class time to cover all of it. (19) Give it a quick read on your own. There are a few slides at the end of today’s notes. Surfactants = surface-active agents e.g., detergents, emulsifiers = amphiphilic substances i.e., have distinct polar & nonpolar regions 1.) ↓ surface tension of water (make water’s skin harder to float on) 2.) act as emulsifying agents (allow immiscible substances to mix)  Hydrophobic tails HYDROPHILIC POLAR HEAD pack together Interacts well with water  form water-soluble globules: “micelles” HYDROPHOBIC NONPOLAR TAIL Interacts well with fats/oils/greases…  Nonpolar cores of & “tails” of other detergent molecules micelles can host oils/greases… Detergent micelle Kotz Figure 13.22 (20)water, salad dressing…) (soapy Emulsions = colloidal suspensions, not solutions. Stable emulsions are colloidal suspensions, not solutions Mix oil & water, Emulsion plus emulsifier… = suspension of micelles… Kotz Figure 13.22 Hydrophilic outside Hydrophobic inside Interacts well with H2Os NOT in contact with H2Os Dissolved in the water NOT dissolved in the water! Optically: Optically: homogeneous very different refractive index looks similar to water  scatters light Make your own stable emulsion: pernod + water = pastis (instantly…) colourless colourless milky… (21) Amphiphilic substances are great emulsifiers e.g., lecithin: a phospholipid in eggs, mustard… Fat-soluble fatty acid tails (varied structures) HH H H HH HH HH HH CH3 HH HH HH O CH3 C C C C N C C C C C C C C C O O C C C CH3 C C H H HH HH O O P HH HH HH HH HH HH C C O HH HH O C H HH H HH HH C C C C HH HH HH C C C C O H HH HH C C HH C C HH C C C C HH HH C C HH HH H H Stabilizes oil-vinegar “emulsions” (oil droplets suspended in water)  essential for making a vinaigrette salad dressing stay mixed Interested in the science involved in cooking? Check out: (22) On Food and Cooking: the Science and Lore of the Kitchen, by Harold McGee.  How much solute is dissolved? What conditions influence solubility? (23) How much solute is dissolved? QUANTITATIVE… SS 13.1 Units of concentration: quantifying solute Fractional composition: sum of parts = whole MASS PERCENT = mA × 100% = mass of solute A (wt %) mA+ mB + mC + … per total mass soln mA mB mC … + + m =1 Often used commercially mTOTAL mTOTAL TOTAL e.g., vinegar = 5% acetic acid 95% water MOLE FRACTION = nA = moles of solute A (χA) nA+ nB + nC + … per total moles soln χA + χB + χC + … = 1 Useful soon for predicting solutions’ vapour pressures. (24) Other important concentration units: SS MOLARITY = mol / L = moles of solute per litre of solution (M) (default unit for chemists) MOLALITY = mol / kg = moles of solute per kg of solvent (m ) * mass is unaffected by temperature * (useful when analyzing melting & boiling behaviour of solutions) PARTS PER MILLION = mg / kg = mg of solute per kg solution (ppm) (used for dilute solutions, i.e.,. mass units of solute. e.g., natural & bottled water) 106 mass units of solution also: parts per billion (ppb) 1 ppb = 1 µg solute/kg soln (25) We will see these in problems. Practice on your own first!  Example: mass percent & molality… & freezing pt? The electrolyte in car batteries (12 V lead storage batteries) is a 3.75 M sulfuric acid solution with a density of 1.230 g/mL. Calculate the mass percent and molality of the sulfuric acid. Hint: solvent is water… H2SO4 MM=98.07 g/mol ANS: Molality (nsolute/kgsolvent) is linked to density and molarity…start there. 1 L has mass of 1230 g = 1.230 kg total = msolute + msolvent and 3.75 M means 1 L contains 3.75 mol H2SO4 = 367.8 g H2SO4 thus msolvent (in 1L soln) = 1.230 – 0.3678 kg = 0.8623 kg solvent  Molality = 3.75 mol H2SO4 = 4.35 mol/kg = 4.35 m 0.8623 kg H2O Mass %: 100% x msolute = 367.8g H2SO4 x100% = 29.9% (wt) H2SO4 msolution 1230g total What is the battery acid’s freezing point? Next class… (…car won’t start if battery is frozen…) (26) How much solute is dissolved? Qualititative… Concentration = amount solute dissolved per amount solvent +? Dilute Concentrated SATURATED  Eq’m system: equal rates of  Unsaturated solutions: dissolving & precipitating conc. < sat’d  more solute can dissolve  Conc’n: max amount solute for that solvent at that T  Solubility = conc. solute dissolved in saturated sol’n (e.g., g solute/100mL water) (27) Supersaturated Solutions (see Closer Look, Kotz 9th Ed. p.474)  metastable: too much solute dissolved  non-equilibrium system, changing slowly  any disturbance  solute precipitates suddenly A very supersaturated solution of sodium acetate crystallizing  (to form a solid salt with water trapped in its crystal lattice)  Crystallization of a supersaturated solution is often accompanied by heat transfer Also possible (and metastable…):  superheated liquids - which may abruptly boil (“bump”)  supercooled liquids - which may abruptly freeze (see in Expt #3?)  supercooled vapours - which may abruptly condense (28) 13.3 Environmental factors affecting solubility: SS pressure (P) & temperature (T- already done) Partial pressure of dissolving gas above solution: determines frequency of collisions between gas & solution surface  determines [gas] at which have balance of rates of gas dissolving & leaving solution (equilibrium!) = a saturated sol’n, where [gas] = “solubility” Table 13.2 Henry’s law: C=kH P or Sg=kH Pg concentration of dissolved gas IS DIRECTLY PROPORTIONAL TO partial P of that gas above the solution (provided gas does not react with solvent!) IMPORTANT: ** unit analysis crucial here, especially P units! ** often given solubility data at given partial P of gas (calculate own kH) (29) Zumdahl’s GASES DISSOLVE BECAUSE MOLECULES SS Figure 11.5 COLLIDE WITH THE LIQUID’S SURFACE… SOLUTION AT CHANGE: RESULT: EQUILIBRIUM ↑ PRESSURE ↑ [dissolved gas] Balance between  ↑ collisions gas molecules with surface  ↑ rate of escape entering (dissolving)  ↑ rate of  new equilibrium & escaping from dissolution reached solution (30)  Example: supersaturated solutions containing GAS SS Sudden ↓ in pressure  gas erupts from solution! Tragedy at Lake Nyos, Cameroon (Kotz 9th p.482) Lake in deep volcanic crater  active vents beneath, emitting CO2(g) Poor turnover of water  depths sat’d with CO2(g) at high P CO2 Aug. 21, 1986: Tremor? ∆ in temperature?  sudden turnover of water in lake  high-[CO2] water now seeing low P CO2 of air 260-foot geyser of CO2 dense gas cloud hugged ground; 45 mph 1700 people died, 12 miles away Now: PREVENTION: 200 m pipes vent depths (31) ON OWN: Why does soda pop go flat? SS Scenario: A can of carbonated soft drink is removed from the refrigerator, opened and left to stand a while…and it “goes flat”. Science: Which choice best describes the changes that cause the soda to lose its CO2 content? A. Salt content was increased B. Temperature was increased C. N2(g) pressure over the solution was raised D. CO2(g) pressure over the solution was lowered E. …more than one of these choices Supersaturated solutions of gas can slowly or abruptly release gas…  nucleation sites speed up approach to equilibrium (gas release!)  Mentos & Coke: http://www.newscientist.com/article/dn14114-science-of-mentosdiet-coke-explosions-explained.html#.U5dzHnY8N8E http://www.youtube.com/watch?v=VlA-zkZssLs (32) Example: CO2 in your soda pop (from an old exam) SS The solubility of gaseous CO2 in water at 10°C is 0.240 g per 100.0 mL, under a pressure of 1.0 atm of carbon dioxide. During manufacturing, a soft drink at 10.0°C is saturated with CO2 under a pressure (P CO2) of 4.0 atm, and then sealed. a) What mass of CO2 is dissolved in a 355mL can of this drink? Ans: k = 0.0024 g/(mL⋅atm) mCO2 = 3.4 g dissolved b) Imagine you open a can of this beverage and leave it open to the atmosphere (PTOTAL = 1.0 atm) at 10.0°C to go “flat”. [Air is ~ 0.031% CO2 by mole; thus, PCO2 =...] What volume of CO2 will be released from the beverage? Ans: mCO2-dissolved = 0.000264 g mCO2-released = 3.4077 g (all…) VCO2 at 1 atm = 1.8 L (33)  ASSIGNED READINGS: BEFORE NEXT CLASS: Read: Ch.11 (all) plus Ch.13.1-13.3 WORK ON: Ch.11,13 problems & conc. unit problems especially Kotz’s end-of-chapter General Questions, In the Laboratory & Summary and Conceptual Questions Review: molarity (Ch.3), mass % (Ch.1&2) practice interconverting concentration units Next week: finish Ch.13… (34) Extra slides: Examples to practice with (35) CLICKER Q: Increasing a substance’s solubility You have a saturated solution of NaCl at 25°C: NaCl(s) → NaCl(aq) No solid is present in the beaker. ∆H°soln[NaCl] = +2.75 kJ/mol ∆S°soln[NaCl] = +43.4 J/mol⋅K How could you increase the amount of dissolved NaCl in this solution? A. Add some more solid NaCl. B. Lower the temperature. C. Raise the temperature. D. Lower the temperature & add some NaCl. E. Raise the temperature & add some NaCl. (36) Extra concentration unit question: mass %, mass fraction Vinegar is a 3-5% (wt %) solution of acetic acid in water. How can we prepare such a solution starting with what chemical companies sell, i.e., glacial acetic acid (17.4 M)? (Acetic acid MM=60.05 g/mol; water MM=18.02 g/mol) And once it’s made: what is its concentration in molarity? ANS: arbitrarily decide to make 100g of it (since not specified) To prepare a 5% (wt) solution: carefully add 5 g of glacial acetic acid to 95 g of H2O, & mix… To calculate molarity: c = nsolute/Vsoln Need to find: #moles acetic acid AND total volume of solution 1.) #mol acetic acid in 100g soln = (5g) / (60.05g/mol) = 0.0833 mol 2.) To find Vsoln: need density (don’t have); so, make an assumption  assume solution has same density as water (1.00g/mL)  then can convert total mass of 100g to volume = 100mL (37) THUS: molarity = (0.0833 mol solute) / (0.1L) = 0.8 M 1 SF Extra slides: colloids (13.5) For your information… (38) Colloidal Dispersions (or, “colloids”) = stable suspensions of tiny particles in a dispersing medium or very high MM molecules (proteins, starches) dissolved in medium Table 13.5 Don’t get hung up on details here… familiar Particles are large enough to see, but too small to settle out Two general kinds of colloids (based on dispersed phase): 1.Hydrophilic colloids: strong attractions to water (e.g., H-bonds) (39) 2.Hydrophobic colloids: not…more detail soon. Why do colloidal particles stay separated (∴ suspended) ? SHOWN HERE: A hydrophobic colloid  Forces between particles dispersing medium & too weak for dissolution e.g., highly insoluble salts finely divided metals soil particles Tiny particle (1-1000 nm) of suspended material at core Kotz Fig.14.21  Colloidal particles evidently “coated” with dissolved ions (even if they aren’t charged themselves…quite interesting)  Outer surfaces all same charge  ELECTROSTATIC REPULSION (40)  To “break” a colloidal suspension REMOVE THE SURFACE CHARGES ON PARTICLES!  ONE WAY = Add an electrolyte ( ↑ dissolved ions…) surface ions attracted elsewhere, not just to particles removes surface charge  no longer repel each other RESULT: pptn of suspended material = “Coagulation” What if you can’t add ions? (…because not dealing with a liquid…) REMOVING SOOT FROM SMOKE (= colloidal dispersion of dust in air) Zumdahl’s Figure 11.25  charged plates attract colloidal particles  impact knocks ions off particles  particles aggregate & ppt as solid soot