Ions in Aqueous Solutions and Colligative Properties

1. Ions in Aqueous Solutions and Colligative Properties Chapter 14

2. Dissociation • The separation of ions that occurs when an ionic compound dissolves.

3. Dissociation examples

4. You try • Write the equation for the dissolution of NH4NO3 in water. If 1 mol of ammonium nitrate is dissolved, how many moles of each type of ion are produced? • 1 mol of each type of ion

5. Precipitation Reactions • When two solutions are mixed, a double replacement reaction may occur. • If one of the products is insoluble, it will form a precipitate. • See table 14-1 on page 427

6. Example • Solutions of (NH4)2S and Cd(NO3)2 are mixed. Will a precipitate form?

7. Example continued The cadmium sulfide is the precipitate.

8. Net Ionic Equations • Includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution.

9. Spectator Ions • Ions that do not take part in a chemical reaction and are found in solution both before and after the reaction.

10. Example

11. You try • A solution of sodium sulfide is combined with a solution of iron(II) nitrate. Does a precipitate form? • Iron(II) sulfide is the precipitate.

12. You try continued • Write the net ionic equation for the previous reaction.

13. Ionization • The process that forms ions from solute molecules by the action of the solvent. • The attraction between the solvent and the solute is strong enough to break the covalent bonds.

14. Hydronium • H3O+ • Formed when an H+ ion is combined with a water molecule (hydrated). • Happens instantly when H+ ions are in water. • Highly exothermic • Formed by many molecular compounds that ionize

15. A more accurate picture

16. Strong electrolytes • Any compound whose dilute aqueous solutions conduct electricity well. • All or almost all dissolved compound is in the form of ions • Not all compound has to dissolve, but the part that does must be ions

17. Weak electrolytes • Any compound whose dilute aqueous solutions conduct electricity poorly. • A small amount of the dissolved compound is in the form of ions.

18. Be careful! • Strong electrolytes have a high degree of ionization or dissociation, regardless of their concentration. • Weak electrolytes have a low degree of ionization or dissociation, regardless of their concentration.

19. Colligative Properties • Properties of solutions that depend on the concentration of solute particles, but not the identity of solute particles.

20. Nonvolatile substance • Has little tendency to become a gas under existing conditions.

21. Vapor-pressure lowering • The vapor pressure of a solvent containing a nonvolatile solute is lower than the vapor pressure of the pure solvent at the same temperature. • The solute lowers the concentration of solvent molecules at the surface. • Fewer molecules enter the vapor phase.

22. Effects • See figure 14-6 on page 436 • The solution remains liquid over a wider temperature range. • The freezing point is lowered and the boiling point is raised.

23. Molal freezing-point constant • Kf • The freezing point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute. • = -1.86 °C/m for water • 2 molal decreases 3.72 °C

24. Freezing-point depression • The difference between the freezing points of the pure solvent and a solution of a nonelectrolyte in that solvent. • It is directly proportional to the molal concentration of the solution.

25. Example • Determine the freezing point of a water solution of fructose, C6H12O6 made by dissolving 58.0 g of fructose in 185 g of water. • -3.24 °C

26. You try • Determine the molal concentration of a solution of ethylene glycol, HOCH2CH2OH, if the solution’s freezing point is -6.40 °C. • 3.44 m

27. Molal boiling-point constant • The boiling-point elevation of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute. • Kb = 0.51 °C/m for water

28. Boiling-point elevation • The difference between the boiling points of the pure solvent and a nonelectrolyte solution of that solvent. • Directly proportional to the molal concentration of the solution

29. Example • What is the boiling point of a solution of 25.0 g of 2-butoxyethanol, HOCH2CH2OC4H9, in 68.7 g of ether? • 40.8 °C

30. You try • What mass of glycerol, CH2OHCHOHCH2OH, must be dissolved in 1.00 kg of water in order to have a boiling point of 104.5 °C? • 810 g

31. Semipermeable membrane • Allows the movement of some particles while blocking the movement of others. • Example: allows water molecules through, but not sucrose molecules

33. Osmosis • The movement of solvent through a semipermeable membrane from the side of lower solute concentration to the side of higher solute concentration.

34. Osmotic pressure • The external pressure that must be applied to stop osmosis. • The greater the concentration of a solution, the greater the osmotic pressure.

35. Electrolytes • 1 mole of an electrolyte produces more than one mole of particles in solution. • The ions separate

36. Example • A water solution contains 42.9 g of calcium nitrate dissolved in 500. g of water. Calculate the freezing point of the solution. • -2.92 °C

37. You try • What is the expected boiling point of a 1.70 m solution of sodium sulfate in water? • 102.6 °C

38. Actual values • Our expected values are not always what is observed. • See table 14-3 on page 445 • Differences are caused by attractive forces between ions in solution.