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Chapter 13 Properties of Solutions

Chapter 13 Properties of Solutions. Solutions. Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent . Solutions.

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Chapter 13 Properties of Solutions

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  1. Chapter 13Properties of Solutions

  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. 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.

  6. 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

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

  8. Breaking attractive interactions is always a _________________ process. • Forming attractive interactions is always a ___________________ process.

  9. 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.

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

  11. 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.

  12. 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.

  13. In the process illustrated below, water vapor reacts with excess sodium sulfate to form the hydrated form of the salt. The chemical reaction is: Na2SO4(s) + 10 H2O  Na2SO410 H2O • Essentially all the water vapor in the closed container is consumed in this reaction. If we consider our system to consist initially of the solid sodium sulfate and water vapor (a) does the system become more or less ordered in the process, and (b) does entropy of the system increase or decrease?

  14. Does the entropy of the system increase or decrease when the stopcock is opened to allow the mixing of the two gases in the apparatus?

  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. • The above reaction is nickel metal and HCl • The residue left after evaporation is chemically different from either reactant, so a chemical reaction took place

  17. 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.

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

  19. 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.

  20. 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.

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

  22. Sample 13.2 pg. 539 • Consider C7H16, Na2SO4, HCl, and I2 individually • Which is the best solvent for each, CCl4 or water?

  23. Arrange the following in order of increasing solubility in water

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

  25. 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.

  26. 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.

  27. Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 of 4.0 atm over the liquid at 25 ⁰ C. The Henry’s law constant for CO2 in water at this temperature is 3.1 x 10-2 mol/ L - atm

  28. Calculate the concentration of CO2 in a sode drink after the bottle is opened and equilibrates at 25 ⁰ C under a CO2 partial pressure of 3.0 x 10-4 atm.

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

  30. 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.

  31. Ways of Expressing Concentrations of Solutions

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

  33. 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 =

  34. moles of A total moles in solution XA = Mole Fraction (X) • In some applications, one needs the mole fraction of solvent, not solute—make sure you find the quantity you need!

  35. mol of solute L of solution M = Molarity (M) • Probably the most familiar concentration unit • Because volume is temperature dependent, molarity can change with temperature.

  36. mol of solute kg of solvent m = Molality (m) Because both moles and mass do not change with temperature, molality (unlike molarity) is not temperature dependent.

  37. A solution is made by dissolving 13.5 g of glucose (C6 H12 O6 ) in 0.100 kg of water. What is the mass percentage of solute in this solution?

  38. A 2.5 gram sample of groundwater is found to contain 5.4 μg of Zn2+ . What is the concentration of Zn2+ in parts per million?

  39. Calculate the mass percentage of NaCl in a solution containing 1.50 grams of NaCl in 50.0 g of water.

  40. A commerical bleaching solution contains 3.62 mass % sodium hypolchlorite, NaOCl. What is the mass of NaOCl in a bottle containing 2500 g of bleaching solution?

  41. A solution is made by dissolving 4.35 grams of glucose in 25.0 mL of water at 25 ⁰ C. Calculate the molality of glucose in the solution.

  42. What is the molality of a solution made by dissolving 36.5 g of naphthalene (C10H8) in 425 g of toluene (C7H8)?

  43. An aqueous solution of HCl contains 36% HCl by mass. Calculate the mole fraction of HCl in the solution. • Calculate the molality of HCl in the solution.

  44. Try on your own • A commercial bleach solution contains 3.62% mass NaOCl in water. Calculate (a) the molality and (b) the mole fraction of NaOCl in the solution.

  45. Changing Molarity to Molality If we know the density of the solution, we can calculate the molality from the molarity, and vice versa.

  46. A solution contains 5.0 grams of toluene (C7H8) and 225 g of benzene and has a density of 0.876 g/mL. Calculate the molarity of the solution.

  47. A solution containing equal masses of glycerol (C3H8O3) and water has a density of 1.10 g/mL. Calculate (a) the molality of glycerol, (b) the mole fraction of glycerol, (c) the molarity of glycerol in the solution.

  48. Colligative Properties • Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. • Among colligative properties are • Vapor pressure lowering • Boiling point elevation • Melting point depression / freezing point depression • Osmotic pressure

  49. Vapor Pressure Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.

  50. Vapor Pressure Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

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