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Chapter 13 Properties of Solutions

Chapter 13 Properties of Solutions

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Chapter 13 Properties of Solutions

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  1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 13Properties of Solutions John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

  2. Solutions • Solutions are homogeneous mixtures of two or more pure substances. • In a solution, the solute is dispersed uniformly throughout the solvent.

  3. Solutions The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

  4. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

  5. Dissolution video

  6. How Does a Solution Form If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

  7. Energy Changes in Solution • Simply put, three processes affect the energetics of the process: • Separation of solute particles • Separation of solvent particles • New interactions between solute and solvent

  8. Energy Changes in Solution The enthalpy change of the overall process depends on H for each of these steps.

  9. Enthalpy activity

  10. Why Do Endothermic Processes Occur? Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered.

  11. Why Do Endothermic Processes Occur? Yet we know that in some processes, like the dissolution of NH4NO3 in water, heat is absorbed, not released.

  12. Energy Changes & Solution Formation • The overall enthalpy change is as follows: DHsoln = DH1 + DH2 + DH3 • If DHsoln > 0, the process is endothermic • If DHsoln < 0, the process is exothermic

  13. Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system.

  14. Enthalpy Is Only Part of the Picture So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.

  15. Student, Beware! Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.

  16. Student, Beware! • Dissolution is a physical change—you can get back the original solute by evaporating the solvent. • If you can’t, the substance didn’t dissolve, it reacted.

  17. Yes, because any addition of solid to liquid significantly changes the entropy. • Yes, because of the energy required to make the AgCl dissolve. • No, because the AgCl is not dispersed throughout the liquid phase.

  18. Yes, because any addition of solid to liquid significantly changes the entropy. • Yes, because of the energy required to make the AgCl dissolve. • No, because the AgCl is not dispersed throughout the liquid phase.

  19. Types of Solutions • Saturated • Solvent holds as much solute as is possible at that temperature. • Dissolved solute is in dynamic equilibrium with solid solute particles.

  20. Types of Solutions • Unsaturated • Less than the maximum amount of solute for that temperature is dissolved in the solvent.

  21. Types of Solutions • Supersaturated • Solvent holds more solute than is normally possible at that temperature. • These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

  22. Yes • No

  23. Yes • No

  24. Factors Affecting Solubility • Chemists use the axiom “like dissolves like”: • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents.

  25. Factors Affecting Solubility The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

  26. Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.

  27. Factors Affecting Solubility • Vitamin A is soluble in nonpolar compounds (like fats). • Vitamin C is soluble in water.

  28. The solubility would be about the same in water as the solubility of glucose. • The solubility would be lower in water than the solubility of glucose. • The solubility would be higher in water than the solubility of glucose.

  29. The solubility would be about the same in water as the solubility of glucose. • The solubility would be lower in water than the solubility of glucose. • The solubility would be higher in water than the solubility of glucose.

  30. SAMPLE EXERCISE 13.2 Predicting Solubility Patterns Predict whether each of the following substances is more likely to dissolve in carbon tetrachloride (CCl4) or in water: C7H16, Na2SO4, HCl, and I2. Solution Analyze: We are given two solvents, one that is nonpolar (CCl4) and the other that is polar (H2O), and asked to determine which will be the best solvent for each solute listed. Plan: By examining the formulas of the solutes, we can predict whether they are ionic or molecular. For those that are molecular, we can predict whether they are polar or nonpolar. We can then apply the idea that the nonpolar solvent will be best for the nonpolar solutes, whereas the polar solvent will be best for the ionic and polar solutes.

  31. SAMPLE EXERCISE 13.2 Predicting Solubility Patterns Predict whether each of the following substances is more likely to dissolve in carbon tetrachloride (CCl4) or in water: C7H16, Na2SO4, HCl, and I2. Solution Solve: C7H16 is a hydrocarbon, so it is molecular and nonpolar. Na2SO4, a compound containing a metal and nonmetals, is ionic; HCl, a diatomic molecule containing two nonmetals that differ in electronegativity, is polar; and I2, a diatomic molecule with atoms of equal electronegativity, is nonpolar. We would therefore predict that C7H16 and I2 would be more soluble in the nonpolar CCl4 than in polar H2O, whereas water would be the better solvent for Na2SO4 and HCl.

  32. SAMPLE EXERCISE 13.2continued PRACTICE EXERCISE Arrange the following substances in order of increasing solubility in water: A B C D Answer:  C5H12 < C5H11 Cl < C5H11 OH < C5H10(OH)2 (in order of increasing polarity and hydrogen-bonding ability). A < D < C < B

  33. From weakest to strongest, rank the following solutions in terms of solvent–solute interactions: NaCl in water, butane (C4H10) in benzene (C6H6), water in ethanol. NaCl in water < C4H10 in C6H6 < water in ethanol Water in ethanol < NaCl in water < C4H10 in C6H6 C4H10 in C6H6 < water in ethanol < NaCl in water

  34. Correct Answer: NaCl in water < C4H10 in C6H6 < water in ethanol Water in ethanol < NaCl in water < C4H10 in C6H6 C4H10 in C6H6 < water in ethanol < NaCl in water Butane in benzene will have only weak dispersion force interactions. Water in ethanol will exhibit much stronger hydrogen-bonding interactions. However, NaCl in water will show ion–dipole interactions because NaCl will dissolve into ions.

  35. Gases in Solution • In general, the solubility of gases in water increases with increasing mass. • Larger molecules have stronger dispersion forces.

  36. Gases in Solution • The solubility of liquids and solids does not change appreciably with pressure. • The solubility of a gas in a liquid is directly proportional to its pressure.

  37. Henry’s Law Sg = kPg where • Sg is the solubility of the gas; • k is the Henry’s law constant for that gas in that solvent; • Pg is the partial pressure of the gas above the liquid.

  38. Henry’s Law Movie

  39. At a certain temperature, the Henry’s law constant for N2 is 6.0  104M / atm. If N2 is present at 3.0 atm, what is the solubility of N2? 6.0  104M 1.8  103M 2.0  104M 5.0  105M

  40. Correct Answer: 6.0  104M 1.8  103M 2.0  104M 5.0  105M Henry’s law, Sg = (6.0  104M/atm)(3.0 atm) Sg = 1.8  103M

  41. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

  42. Temperature • The opposite is true of gases: • Carbonated soft drinks are more “bubbly” if stored in the refrigerator. • Warm lakes have less O2 dissolved in them than cool lakes.

  43. Gases are emitted from the cooking pot surfaces as it is heated. • Dissolved gases are less soluble in solution as temperature increases. • Water molecules begin to enter the gas phase to stimulate boiling. • Boiling actually begins on a small scale at temperatures below the boiling point.

  44. Gases are emitted from the cooking pot surfaces as it is heated. • Dissolved gases are less soluble in solution as temperature increases. • Water molecules begin to enter the gas phase to stimulate boiling. • Boiling actually begins on a small scale at temperatures below the boiling point.

  45. Ways of Expressing Concentrations of Solutions

  46. mass of A in solution total mass of solution Mass Percentage Mass % of A =  100

  47. Determine the mass percentage of hexane in a solution containing 11 g of butane in 110 g of hexane. 9.0 % 10. % 90.% 91 %

  48. Correct Answer: 9.0 % 10. % 90.% 91 % mass of component in solution =  mass % of component 100 total mass of solution Thus, 110 g (110 g + 11 g)  100 = 91%

  49. mass of A in solution total mass of solution mass of A in solution total mass of solution Parts per Million andParts per Billion Parts per Million (ppm) ppm =  106 Parts per Billion (ppb)  109 ppb =

  50. .23 ppm • 2.3 ppm • 230 ppm • 2300 ppm