1 / 44

Colligative Properties

Colligative Properties. How does the solute change the properties of the solvent?. Consider aqueous solutions. Solvent = water. How do the properties of the solution compare to the properties of pure water?. Conductivity. Pure water does not conduct electricity. Some solutions do.

teagues
Télécharger la présentation

Colligative Properties

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Colligative Properties

  2. How does the solute change the properties of the solvent? • Consider aqueous solutions. • Solvent = water. • How do the properties of the solution compare to the properties of pure water?

  3. Conductivity • Pure water does not conduct electricity. • Some solutions do. • In order to conduct a current, a solution must contain ions. • Conductivity depends on the nature of the particle – Molecular solutes do not conduct.

  4. Colligative Properties • Depend on the number of particles in solution not on the type of particles. • Doesn’t matter if particles are ions or molecules. • Concentration of particles doesmatter.

  5. What are some colligative properties? • Vapor Pressure Lowering • Freezing Point Depression • Boiling Point Elevation • Osmotic Pressure • The higher the concentration of solute particles, the larger the effect!

  6. Colligative Properties • Useful for determining the nature of solute after it is dissolved in the solvent • Useful for determining the molecular mass of the solute

  7. Vapor Pressure Lowering • Presence of nonvolatile solute lowers the vapor pressure of the solution compared to that of the pure solvent. • Has implications for the phase diagram. • Normal b.p. = temperature at which the v.p. = 1 atm. • So, solutions must be heated to a higher T to reach a v.p. of 1 atm. • Boiling Point Elevation

  8. Nonvolatile Solute Pure solvent Solvent + a nonvolatile solute (nrp) New triple point.

  9. Change in boiling point, Tb • Tb = Tb – Tb0 • Tb = boiling point of solution – boiling point of solvent (Tb > Tb0) so Tb is always a positive #. Tb = Kbmsolute • Kb = a constant characteristic of the solvent = molal boiling-point elevation constant (C/m) • msolute = molality of the solute in solution

  10. Tb = Kbmsolute • The higher the concentration, the bigger the effect. So solutes stretch out the temperature range of the liquid phase. More concentrated solutions  100C 0C Pure H2O

  11. Boiling Point of Solution • What is the boiling point of a solution containing 478 g of ethylene glycol (antifreeze) in 3202 g of water? • The molar mass of ethylene glycol is 62 g/mol. • Kb for water is 0.51Ckg/mol. • Boiling point of pure water = 100C

  12. Boiling Point of Solution • Tb = Kbmsolute • Moles ethylene glycol = 478g / 62 g/mol =7.71 • Molality ethylene glycol = 7.71 mol / 3.202 kg H2O • = 2.41 m • Tb = (0.51Ckg/mol)(2.41 mol/kg) = 1.23C • Boiling Point of Solution = Tb + Tb = 101.23C

  13. Calculating Molar Mass • msolute = Tb/Kb • msolute = moles solute/kg solvent • And moles solute = mass solute/FM •  (mass solute/FM) = Tb/Kb kg solvent

  14. Calculating Molar Mass • Rearranging: • FM = (kg solvent)(Tb) Kb(mass solute)

  15. Kb(mass solute)FM = (kg solvent)(Tb) • Tb of pure benzene = 80.1C • Kb for benzene is 2.53 Ckg/mol • 10.9 grams of unknown solute are dissolved in 75.8 grams of benzene. Tb of the solution is 82.1C. Estimate the molar mass of the solute. • So Tb = 2.0C

  16. Kb(mass solute)FM = (kg solvent)(Tb) • FM = (2.53Ckg/mol)(10.9 g) (0.0758 kg)(2.0C) FM = 182 g/mol

  17. Change in freezing point, Tf • Tf = Tf0 – Tf • Tf = freezing point of solvent – freezing point of solution. (Tf0 > Tf) always positive Tf = Kfmsolute • Kf = molal freezing-point depression constant of the solvent (C/m) • m = molality of solute in solution

  18. What is the freezing point of a solution containing 478 g of ethylene glycol in 3202 g of H2O? • Formula mass of ethylene glycol = 62 g/mol. Moles of ethylene glycol = 478 g/62 g/mol = 7.71 moles. For water, Kf = 1.86 C/m. Molality = 7.71 moles/3.202 kg = 2.41 m. • Tf = Kfm = (1.86C/m)(2.41 m) = 4.48C. • Tf = 0.0C – 4.48C = -4.48C

  19. Does it matter if the solute is an electrolyte or a nonelectrolyte? • The number of particles will vary. • Electrolytes will dissolve to produce positive and negative ions  more particles per mole of solute. • For nonelectrolytes, 1-to-1 relationship between moles of solute and moles of dissolved particles.

  20. C6H12O6 • Nonelectrolyte • Dissolves as molecules • C6H12O6(s) C6H12O6(aq) • 1 mole of sugar yields 1 mole of molecules

  21. NaCl • Electrolyte (salt) • Dissolves as ions • NaCl(s)  Na+(aq) + Cl-(aq) • 1 mole of salt yields 2 moles of ions

  22. MgCl2 • Electrolyte • Dissolves as ions • MgCl2(s)  Mg2+(aq) + 2Cl-(aq) • 1 mole of salt yields 3 moles of ions

  23. Nature of solute • Colligative properties depend on concentration of solute particles • Use experimental data to determine moles of particles in solution.

  24. Real Life • Salts do not always exhibit complete dissociation. • Result: The actual Tb or Tf is less than predicted from the dissolving equation. • Evidence that “Ion-pairing” occurs in solution. • More ion pairing in concentrated solutions & when ions have multiple charges.

  25. As a solute is added to a solvent, what happens to the freezing point & the boiling point of the solution? 1) The freezing point decreases & the boiling point decreases. 2) The freezing point decreases & the boiling point increases. 3) The freezing point increases & the boiling point decreases. 4) The freezing point increases & the boiling point increases.

  26. Which solution containing 1 mole of solute dissolved in 1000 g of water has the lowest freezing point? • C2H5OH(aq) • NaCl(aq) • KOH 4) CaCl2 Be careful! What if the question asked which solution has the highest freezing point? Highest FP! Lowest FP!

  27. Of the following solutions, the one that will freeze at the lowest temperature contains 1 mole of nonvolatile solute dissolved in 1)250 g of solvent 2) 500 g of solvent 3) 750 g of solvent 4) 1000 g of solvent

  28. Which solute, when added to 1000 g of water, will produce a solution with the highest boiling point? • 29 g of NaCl • 58 g of NaCl • 31 g of C2H6O2 4) 62 g of C2H6O2 About 1 mole of NaCl. 2 moles of ions. About 1 mole of C2H6O2. 1 mole of molecules.

  29. Which solution will freeze at the lowest temperature? • 1 g of NaCl dissolved per 100 g of H2O • 1 g of NaCl dissolved per 1000 g of H2O 3) 1 g of C12H22O11 dissolved per 100 g of H2O 4) 1 g of C12H22O11 dissolved per 1000 g of H2O

  30. Which solution will freeze at the lowest temperature? 1)1 g of NaCl dissolved per 100 g of H2O 1 g  58.8 g/mole = 0.017 moles  0.034 moles of ions. 3) 1 g of C12H22O11 dissolved per 100 g of H2O 1 g  232 g/mole = 0.0043 moles of molecules

  31. Effect of NonvolatileSolute • Boiling Point Elevation • Freezing Point Depression • The more particles, the bigger the effect!

  32. Applications of colligative properties • Salting roads in the winter time. • Using salted ice to make ice cream. • Antifreeze in radiator to protect water from overheating or freezing

  33. Osmosis

  34. Osmosis • A solution and a pure solvent are separated by a semipermeable membrane. • Solvent, but NOT solute molecules, can pass through the membrane.

  35. Osmosis • The solvent flows from the region of higher solvent concentration to the region of lower solvent concentration. • Volumesolution and volumesolvent  as a f(t) • At equilibrium, the liquid levels stop changing. • More hydrostatic pressure on the solution than on the solvent.

  36. Hydrostatic pressure due to

  37. Another common set-up

  38. Osmotic Pressure • Minimum hydrostatic pressure that stops the net flow of solvent across the membrane into the solution. • Can be used to characterize solutions and determine molar mass • Small concentration of solute  produces large osmotic pressure

  39. Net movement of solvent into solution over time. Note: the phrases “more and less concentrated solution” relate to the initial and final conditions of the solution in the bulb. The pure solvent is outside the bulb. At equilibrium

  40. Net flow of water across membrane. At equilibrium

  41. Reverse Osmosis

More Related