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Properties of Solutions

Properties of Solutions. Classification of Matter. Solutions are homogeneous mixtures. Solute. A solute is the dissolved substance in a solution. Salt in salt water. Sugar in soda drinks. Carbon dioxide in soda drinks. Solvent. A solvent is the dissolving medium in a solution.

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Properties of Solutions

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  1. Properties of Solutions

  2. Classification of Matter Solutions are homogeneous mixtures

  3. Solute A solute is the dissolved substance in a solution. Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda

  4. Calculations of Solution Concentration Mass percent - the ratio of mass (in grams) of solute to mass (in grams) of solution, expressed as a percent

  5. Calculations of Solution Concentration Mass/volume (m/v) % - the ratio of mass (in grams) of solute to volume of solution (in mL), expressed as a percent

  6. Calculations of Solution Concentration Volume/volume (v/v) % - the ratio of volume (in mL) of solute to volume of solution (in mL), expressed as a percent

  7. Calculations of Solution Concentration Mole fraction – the ratio of moles of solute to total moles of solution

  8. Calculations of Solution Concentration Molarity (M) - the ratio of moles of solute to liters of solution

  9. Calculations of Solution Concentration Normality (N) – moles of equilvalents/Liter of solution

  10. Calculations of Solution Concentration Molality (m) – moles of solute per kilogram of solvent

  11. Dissolving Stuff

  12. “Like Dissolves Like” Nonpolar solutes dissolve best in nonpolar solvents Polar and ionic solutes dissolve best in polar solvents

  13. Heat of Solution The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent.

  14. Steps in Solution Formation H1Expanding the solute Separating the solute into individual components H2Expanding the solvent Overcoming intermolecular forces of the solvent molecules H3Interaction of solute and solvent to form the solution

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

  16. Why do endothermic processes sometimes occur spontaneously? Some processes, like the dissolution of NH4NO3 in water, are spontaneous at room temperature even though heat is absorbed, not released.

  17. Enthalpy Is Only Part of the Picture • Entropy is a measure of: • Dispersal of energy in the system. • Number of microstates (arrangements) in the system. • b. has greater entropy,  is the favored state (more on this in chap 19)

  18. Predicting Solution Formation

  19. Solubility Trends • The solubility of MOST solids increases with temperature. • The rate at which solids dissolve increases with increasing surface area of the solid. • The solubility of gases decreases with increases in temperature. • The solubility of gases increases with the pressure above the solution.

  20. Saturation of Solutions • A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated. • A solution that contains less solute than a saturated solution under existing conditions is unsaturated. • A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.

  21. 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.

  22. Solubility Chart

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

  24. 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.

  25. Temperature • 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.

  26. Henry’s Law The concentration of a dissolved gas in a solution is directly proportional to the pressure of the gas above the solution Applies most accurately for dilute solutions of gases that do not dissociate or react with the solvent Yes  CO2, N2, O2 No  HCl, HI

  27. Therefore… Solids tend to dissolve best when: • Heated • Stirred • Ground into small particles Gases tend to dissolve best when: • The solution is cold • Pressure is high

  28. 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 • Freezing point depression • Osmotic pressure increase

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

  30. Raoult’s Law The presence of a nonvolatile solute lowers the vapor pressure of the solvent. Psolution = Observed Vapor pressure of the solution solvent = Mole fraction of the solvent P0solvent = Vapor pressure of the pure solvent

  31. What is the vapor pressure in mmHg of a solution that contains 155 grams of glucose dissolved in 250. mL of water at 25.00 C? The vapor pressure of pure water at 25.00C is 23.78 mmHg. 22.4 mmHg

  32. Liquid-liquid solutions in which both components are volatile Modified Raoult's Law: P0 is the vapor pressure of the pure solvent PAand PB are the partial pressures

  33. At a given temperature, you have a mixture of benzene (vapor pressure of pure benzene = 745 torr) and toluene (vapor pressure of pure toluene = 290. torr). The mole fraction of benzene in the solution is 0.450. Assuming both substances are volatile, calculate the vapor pressure of the solution. 495 torr

  34. 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.

  35. The van’t Hoff Factor, i Electrolytes may have two, three or more times the effect on boiling point, freezing point, and osmotic pressure, depending on its dissociation.

  36. Dissociation Equations and the Determination of i i = 2 NaCl(s)  Na+(aq) + Cl-(aq) i = 2 AgNO3(s)  Ag+(aq) + NO3-(aq) i = 3 MgCl2(s)  Mg2+(aq) + 2 Cl-(aq) i = 3 Na2SO4(s)  2 Na+(aq) + SO42-(aq) AlCl3(s)  Al3+(aq) + 3 Cl-(aq) i = 4

  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. 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.

  41. Boiling Point Elevation and Freezing Point Depression In both equations, T does not depend on what the solute is, but only on how many particles are dissolved. Tb = Kb  i m Tf = Kf  i m

  42. Boiling Point Elevation Each mole of solute particles raises the boiling point of 1 kilogram of water by 0.51 degrees Celsius. Kb = 0.51 C  kilogram/mol m = molality of the solution i = van’t Hofffactor

  43. Boiling Point Elevation The change in boiling point is proportional to the molality of the solution: Tb = Kb i 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.

  44. If I add 45.0 grams of sodium chloride to 500. grams of water, what will the boiling point be of the resulting solution? Kb(H2O) = 0.512 0C/m 101.57 ˚C

  45. Freezing Point Depression Each mole of solute particles lowers the freezing point of 1 kilogram of water by 1.86 degrees Celsius. Kf = 1.86 C  kilogram/mol m = molality of the solution i = van’t Hofffactor

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

  47. If I add 92.0 grams of sodium chloride to 500. grams of water, what will the freezing point be of the resulting solution? Kf(H2O) = 1.86 0C/m -11.7 ˚ C

  48. Freezing Point Depression and Boiling Point Elevation Constants, C/m

  49. Osmotic Pressure The minimum pressure that stops the osmosis is equal to the osmotic pressure of the solution

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