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Chapter 14: Solutions and Their Properties If you’re not part of the solution,

Chapter 14: Solutions and Their Properties If you’re not part of the solution, you’re part of the precipitate!. Properties of Solutions. Physical Properties of Solutions. Have you ever wondered... Why antifreeze keeps water from freezing? Why salt causes ice to melt?

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Chapter 14: Solutions and Their Properties If you’re not part of the solution,

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  1. Chapter 14: Solutions and Their Properties If you’re not part of the solution, you’re part of the precipitate!

  2. Properties of Solutions

  3. Physical Properties of Solutions • Have you ever wondered... • Why antifreeze keeps water from freezing? • Why salt causes ice to melt? • Why cooks add salt to boiling water? • Why root beer foams only when poured? • What force opposes gravity to allow water to climb up a tree?

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

  5. nonelectrolyte weak electrolyte strong electrolyte An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.

  6. “like dissolves like” Two substances with similar intermolecular forces are likely to be soluble in each other. • non-polar molecules are soluble in non-polar solvents • CCl4 in C6H6 • polar molecules are soluble in polar solvents • C2H5OH in H2O • ionic compounds are more soluble in polar solvents • NaCl in H2O or NH3 (l)

  7. Energy Changes in Solution To determine the enthalpy change, we divide the process into 3 steps. • Separation of solute particles. • Separation of solvent particles to make ‘holes’. • Formation of new interactions between solute and solvent.

  8. Three types of interactions in the solution process: • solvent-solvent interaction • solute-solute interaction • solvent-solute interaction DHsoln = DH1 + DH2 + DH3

  9. Enthalpy Changes in Solution The enthalpy change of the overall process depends on H for each of these steps. Start End Start End

  10. Calculating ⁄Hsolution (enthalpy of solution) KF(s) -> K+(g) + F-(g) ⁄ E = +821 kJ K+(g) + F-(g) + H2O -> K+(aq) + F-(aq) ⁄E=-819 kJ So Net - KF(s) -> K+(aq) + F-(aq) ⁄ Hsolution = +2kJ

  11. Degree of saturation • Saturated solution • Solvent holds as much solute as is possible at that temperature. • Undissolved solid remains in flask. • Dissolved solute is in dynamic equilibrium with solid solute particles.

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

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

  14. Gases in Solution • The solubility of liquids and solids does not change appreciably with pressure. • But, the solubility of a gas in a liquid is directly proportional to its pressure. Increasing pressure above solution forces more gas to dissolve.

  15. Henry’s Law What happens to the solubility of carbon dioxide in a bottle of soda when the pressure is reduced?

  16. low P high P low c high c Pressure and Solubility of Gases The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution (Henry’s law). c is the concentration (M) of the dissolved gas c = kP P is the pressure of the gas over the solution k is a constant (mol/L•atm) that depends only on temperature

  17. Chemistry In Action: The Killer Lake 8/21/86 CO2 Cloud Released 1700 Casualties • Trigger? • earthquake • landslide • strong Winds Lake Nyos, West Africa

  18. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature. Solubility is measured as the mass of solute dissolved in 100 g of solvent at a given temperature

  19. Temperature • The opposite is true of gases. Higher temperature drives gases out of solution. • Carbonated soft drinks are more “bubbly” if stored in the refrigerator. • Warm lakes have less O2 dissolved in them than cool lakes.

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

  21. 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!

  22. mol of solute L of solution M = Molarity (M) • Because volume is temperature dependent, molarity can change with temperature.

  23. mol of solute kg of solvent m = Molality (m) Because neither moles nor mass change with temperature, molality (unlike molarity) is not temperature dependent.

  24. Mass/Mass Moles/Mass Moles/Moles Moles/L

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

  26. Colligative Properties • 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 • Osmotic pressure

  27. Vapor Pressures of Pure Water and a Water Solution The vapor pressure of water over pure water is greater than the vapor pressure of water over an aqueous solution containing a nonvolatile solute. Solute particles take up surface area and lower the vapor pressure

  28. Vapor Pressure As solute molecules are added to a solution, the solvent becomes less volatile (has decreased vapor pressure). Solute-solvent interactions contribute to this effect.

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

  30. Colligative Properties • Lowering Vapor Pressure • Raoult’s Law: • Where: PA = vapor pressure with solute, • PA = vapor pressure without solute (pure solvent), and • A = mole fraction of A (the pure solvent).

  31. Boiling Point Elevation and Freezing Point Depression Solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

  32. Boiling Point Elevation The change in boiling point is proportional to the molality of the solution: Tb = Kb  m where Kb is the molal boiling point elevation constant, a property of the solvent. Tb is added to the normal boiling point of the solvent.

  33. Freezing Point Depression • The change in freezing point can be found similarly: Tf = Kf  m • Here Kf is the molal freezing point depression constant of the solvent. Tf is subtracted from the normal freezing point of the solvent.

  34. Colligative Properties • Boiling-Point Elevation • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, mS: • Decrease in freezing point (Tf) is directly proportional to molality (Kfis the molal freezing-point-depression constant): In both equations, T does not depend on what the solute is, but only on how many particles are dissolved.

  35. Colligative Properties of Electrolytes Because these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) show greater changes than those of nonelectrolytes. e.g. NaCl dissociates to form 2 ion particles; its limiting van’t Hoff factor is 2.

  36. Colligative Properties of Electrolytes However, a 1 M solution of NaCl does not show twice the change in freezing point that a 1 M solution of methanol does. It doesn’t act like there are really 2 particles.

  37. van’t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions.

  38. van’t Hoff Factor Some Na+ and Cl− reassociate as hydrated ion pairs, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

  39. The van’t Hoff Factor • Reassociation is more likely at higher concentration. • Therefore, the number of particles present is concentration dependent.

  40. The van’t Hoff Factor We modify the previous equations by multiplying by the van’t Hoff factor, i Tf = Kf  m i i = 1 for non-elecrtolytes

  41. Osmosis • Semipermeable membranes allow some particles to pass through while blocking others. • In biological systems, most semipermeable membranes (such as cell walls) allow water to pass through, but block solutes.

  42. Osmosis In osmosis, there is net movement of solvent from the area of higher solvent concentration (lowersoluteconcentration) to the are of lower solvent concentration (highersoluteconcentration). Water tries to equalize the concentration on both sides until pressure is too high.

  43. Colligative Properties • Osmosis • Osmotic pressure, , is the pressure required to stop osmosis:

  44. Molar Mass from Colligative Properties We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound.

  45. Properties of Solutions

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