1. University Chemistry Chapter 9: Physical Equilibrium Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. 2 Phase Boundaries in Pure Substances Two different phasesSame T and P

3. 3 Clapeyron equation  G ,2-1 =  G ,2-1 lim   0, this equation becomes dG  = dG  S  dT-V  dP = S  dT-V  dP (S  -S  )dT= (V  -V  )dP

4. 4  V fus and  H fus are + dP/dT is + dP/dT is -  V fus is - and  H fus is +  V fus typically is small. P-T slope is large. Large changes in pressure are required to cause even small changes in melting temperature. CO 2 H2OH2O

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7. 7 The Clausius-Clapeyron Equation For a phase transition in which one phase is a gas, e.g. vaporization :

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9. 9 P and T are known, From such a plot, we can determine  H vap For sublimation, use  H vap instead of  H vap

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11. 11 Represents an average over the temperature interval 15 o C to 80 o C.

12. 12 Solubility is determined by Intermolecular Forces, T and P Let’s focus on three kinds of liquid solutions

13. 13 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. Crystals rapidly form when a seed crystal is added to a supersaturated solution. Crystallization is the process in which dissolved solute comes out of solution and forms crystals

14. 14 Three types of interactions in the solution process: solvent-solvent interaction solute-solute interaction solvent-solute interaction  H soln =  H 1  H 2  H 3 If |  H 3 |  H 1  H 2,  H soln  & so the solution process is favorable.

15. 15 non-polar molecules are soluble in non-polar solvents CCl 4 in C 6 H 6 polar molecules are soluble in polar solvents C 2 H 5 OH in H 2 O ionic compounds are more soluble in polar solvents NaCl in H 2 O or NH 3 (l) “like dissolves like” Two substances with similar intermolecular forces are likely to be soluble in each other. Hydrophilic (water loving) hydrophobic (water fearing)

16. 16 “soluble”, “slightly soluble”, or “insoluble” depending on solubility. When two liquids are completely soluble in each other in all proportions they are said to be miscible. Solvation is the process in which an ion or a molecule is surrounded by solvent molecules arranged in a specific manner. hydration when the solvent is water Alcohols are miscible with water

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18. 18 The Cleansing Action of Soap

19. 19 Temperature and Solubility Solid solubility and temperature Red: solubility increases with increasing temperature Blue: solubility decreases with increasing temperature

20. 20 The enthalpy of solution at infinite dilution  H soln is the enthalpy change associated with the addition of a solute to a pure solvent. The  H sat is the enthalpy change associated with the addition of solute to a saturated solution The stable form of an ionic solid at saturation is often a hydrate. The enthalpy of solution for hydrates is generally positive (endothermic), especially at satruation. For this reason, the solubility of an ionic compound may increase with temperature, even though the  H soln for the anhydrous form is negative (exothermic).  H soln versus  H sat Solubility-vs-T relation can be explained with  H sat using an Eq. similar to the C-C equation.

21. 21 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 KNO 3 contaminated with 10 g NaCl. Fractional crystallization: 1.Dissolve sample in 100 mL of water at 60 0 C 2.Cool solution to 0 0 C 3.All NaCl will stay in solution (s = 34.2g/100g) 4.78 g of PURE KNO 3 will precipitate (s = 12 g/100g). 90 g – 12 g = 78 g

22. 22 Gas Solubility and Temperature O 2 gas solubility and temperature solubility usually decreases with increasing temperature  G =  H - T  S Contributes to thermal pollution, the heating of the environment

23. 23 Effect of Pressure on the Solubility of Gases Henry’s law: the equilibrium concentration of a gas dissolved in a liquid is proportional to the pressure of the gas over the solution. molar concentration (mol L -1 ) Henry’s law constant Gas pressure over solution (bar), For gases that react with water (e.g. NH 3 and CO 2 ), the concentration of the dissolved gas is as predicted, but the total amount of the gas that dissolves is underestimated. Also, O 2 in blood. Hb +4 O 2 = Hb(O 2 ) 4

24. 24 The Killer Lake Lake Nyos, West Africa 8/21/86 CO 2 Cloud Released 1700 Casualties Trigger? earthquake landslide strong Winds

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27. 27 Liquid-Vapor Equilibrium of Ideal Solutions A liquid solution that is a mixture of two components, A and B, is called a binary liquid mixture (or binary solution). Raoult’s law mole fraction of ivapor pressure of i in its pure liquid form intercept slope Linear form

28. 28 P benzene = X benzene P* benzene P toluene = X toluene P * toluene P T = P benzene + P toluene P T = X benzene P* benzene + X toluene P* toluene Ideal Solution Binary mixture of benzene and toluene

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31. 31 Total vapor pressure as a function of the vapor composition, Raoult’s law

32. 32 Pressure-composition diagram: liquid-vapor phase diagram for an ideal binary solution at a fixed temperature T. vapor liquid lever rule:

33. 33 Temperature-composition phase diagram dew-point curve (pressure is fixed)

34. 34 Fractional Distillation Apparatus

35. 35 P T is greater than predicted by Raoults’s law P T is less than predicted by Raoults’s law Force A-B Force A-A Force B-B <& Force A-B Force A-A Force B-B >& Nonideal Solutions

36. 36 Azeotrope: a solution for which the liquid phase and the vapor phase have the same composition. low-boiling azeotrope high-boiling azeotrope

37. 37 Colligative Properties Properties of solution phase equilibria that depend only upon the number of solute molecules, not their type For solutes that are nonelectrolytes (i.e., they do not dissociate into ions in solution), colligative properties typically hold for concentrations less than about 0.1 M. Approximations: (1) the solutions are dilute enough that they can be considered to be ideal (2) the solutes are nonelectrolytes.

38. 38 Vapor-Pressure Lowering Solute is nonvolatile - does not have a measurable vapor pressure The vapor pressure of its solution is always less than that of the pure solvent. decrease in the vapor pressure

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41. 41 Boiling-Point Elevation molal boiling-point elevation constant (or ebullioscopic constant) molality of the solution T b > T* b T* b freezing point of the pure solvent T b freezing point of the solution

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43. 43 Freezing-Point Depression T* f > T f  T f > 0 T* f is the freezing point of the pure solvent T f is the freezing point of the solution molal freezing-point depression constant (or cryoscopic constant)

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46. 46 Osmotic Pressure (  ) Osmosis is the selective passage of solvent molecules through a porous membrane from a dilute solution to a more concentrated one. A semipermeable membrane allows the passage of solvent molecules but blocks the passage of solute molecules. Osmotic pressure (  ) is the pressure required to stop osmosis. dilute more concentrated

47. 47 Red blood cells and Osmosis isotonic solution hypotonic solution hypertonic solution

48. 48 30 bar

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52. 52 Colligative Properties of Electrolyte Solutions Electrolytes: substances that dissociate into ions in solution. van’t Hoff factor An ion pair is made up of one or more cations and one or more anions held together by electrostatic forces. Ion pair formation affects the number of particles in the solution.

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