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 constantKf
The freezing point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute.
= °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 °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 °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 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.