1 / 23

Chapter 13 Properties of Solutions

Chapter 13 Properties of Solutions. Consider KCl (solute) dissolving in water (solvent): H-bonds in water have to be interrupted, KCl dissociates into K + and Cl - , ion-dipole forces form: K + …  -OH 2 and Cl - …  +H 2 O. ions are solvated by water.

nyx
Download Presentation

Chapter 13 Properties of Solutions

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. Chapter 13Properties of Solutions

  2. Consider KCl (solute) dissolving in water (solvent): • H-bonds in water have to be interrupted, • KCl dissociates into K+ and Cl-, • ion-dipole forces form: K+…-OH2 and Cl- …+H2O. • ions are solvated by water.

  3. Energy Changes and Solution Formation • There are three enthalpy steps in forming a solution: • separation of solute molecules (H1), • separation of solvent molecules (H2), andformation of solute-solvent interactions (H3). • The enthalpy change in the solution process is • Hsoln = H1 + H2 + H3. • Hsoln can either exothermic or endothermic depending on the intermolecular forces.

  4. Energy Changes and Solution Formation • To determine whether Hsoln is positive or negative, consider the strengths of all solute-solute and solute-solvent interactions: • H1 and H2 are both positive. • H3 is always negative. • It is possible to have either H3 > (H1 + H2) or H3 < (H1 + H2).

  5. Energy Changes and Solution Formation • Examples: • NaOH added to water has Hsoln = -44.48 kJ/mol. • NH4NO3 added to water has Hsoln = + 26.4 kJ/mol. • “Rule”: polar solvents dissolve polar solutes. Non-polar solvents dissolve non-polar solutes. • If Hsoln is too endothermic a solution will not form. • NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions.

  6. Solution Formation, Spontaneity, and Disorder

  7. Saturated Solutions and Solubility • Saturation: crystallization and dissolution of a solute are in equilibrium. • Solubility: amount of solute required to form a saturated solution. • Supersaturated: a solution formed when more solute is dissolved than in a saturated solution.

  8. Factors Affecting Solubility • Polar liquids tend to dissolve in polar solvents. • Miscible liquids: mix in any proportions. • Immiscible liquids: do not mix. • The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water. • The number of -OH groups within a molecule increases solubility in water. • Generalization: “like dissolves like”. • The more polar bonds in the molecule, the better it dissolves in a polar solvent. • The less polar the molecule the less it dissolves in a polar solvent and the better is dissolves in a non-polar solvent.

  9. Solute-Solvent Interaction

  10. Solute-Solvent Interaction • Network solids do not dissolve because the strong intermolecular forces in the solid are not re-established in any solution. • Pressure Effects • Solubility of a gas in a liquid is a function of the pressure of the gas.

  11. Pressure Effects • The higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution. • Therefore, the higher the pressure, the greater the solubility. • The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility. • Carbonated beverages are bottled with a partial pressure of CO2 > 1 atm. As the bottle is opened, the partial pressure of CO2 decreases and the solubility of CO2 decreases. Therefore, bubbles of CO2 escape from solution.

  12. Temperature Effects • Experience tells us that sugar dissolves better in warm water than cold. • As temperature increases, solubility of solids generally increases. • Sometimes, solubility decreases as temperature increases (e.g. Ce2(SO4)3).

  13. Temperature Effects • Experience tells us that carbonated beverages go flat as they get warm. • Therefore, gases get less soluble as temperature increases. • Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble and are not available for plants or animals.

  14. Units of Concentration – Interconversion Chart.

  15. Colligative Properties • Colligative properties depend on quantity of solute molecules. (E.g. freezing point depression and melting point elevation.)

  16. Lowering Vapor Pressure • Non-volatile solvents reduce the ability of the surface solvent molecules to escape the liquid. • Therefore, vapor pressure is lowered. • The amount of vapor pressure lowering depends on the amount of solute.

  17. Lowering Vapor Pressure • Raoult’s Law: The vapor pressure of an ideal solution (PA) is a fraction of the vapor pressure of the pure solvent (PA).

  18. Boiling-Point Elevation • At 1 atm (normal boiling point of pure liquid) there is a lower vapor pressure of the solution. Therefore, a higher temperature is required to teach a vapor pressure of 1 atm for the solution (Tb). • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with solute molality, m:

  19. Freezing Point Depression • The solution freezes at a lower temperature (Tf) than the pure solvent. • Decrease in freezing point (Tf) is directly proportional to solute molality (Kfis the molal freezing-point-depression constant):

More Related