1. Chapter 14: Ions in Aqueous Solutions and Colligative Properties Day 1: Dissociation, Solubility, Net Ionic Equations, Ionization Day 2: Colligative Properties, Lab =) Day 3: Review..study guide will be graded today Day 4: Test

2.  Day 1: Dissociation, Ionization, Solubility, Net Ionic Equations

3. Dissociation  Dissocation is when an ionic compound dissolves in water and the ions separate from each other.  Examples:  NaCl (s)  Na +1 (aq) + Cl -1 (aq)  CaCl 2 (s)  Ca +2 (aq) + 2 Cl -1 (aq)  Note the moles!  0.22 moles CaCl 2  0.22 moles of Ca +2 + 0.44 moles Cl -1

4. Example  If 0.484 moles barium nitrate dissociates, how many moles of each ion or polyatomic ion is formed?  Write dissociation equation:  Finish: Barium nitrate and Aluminum

5. Ionization  Some polar molecules can form ions in solution…ions are formed where none existed in the undissolved compound.  HF + H 2 0  H 3 0 + + F -  CH 3 COOH + H 2 0  CH 3 COO - + H 3 0 +

6. These dissociate in water: 1.Strong acids: 2.Strong bases: 3.Soluble salts: Therefore by deduction…don’t dissociate anything else…just leave it “together”.

7. Copyright © Houghton Mifflin Company. All rights reserved. 4b–7

8. Write dissociation equations for: 1.Sulfuric acid 2.Calcium hydroxide 3.Potassium nitrate 4.Sodium oxalate 5.Ammonium chloride 6.Hydrochloric acid 7.Perchloric acid

9. Net Ionic Equations  These contain only those compounds or ions that undergo a chemical change in a reaction in aqueous solution ( net ionic equations contain no spectators!).

10. Example: Consider the rxn of ammonium sulfide and cadmium (II) nitrate. Write molecular, complete ionic and net ionic reactions.  Molecular:  Complete ionic  Net ionic reaction

11. Example: Reaction of hydrochloric acid and sodium hydroxide.  Molecular:  Complete ionic  Net ionic reaction

12. Example: Acetic acid and sodium hydroxide  Molecular:  Complete ionic  Net ionic reaction

13. Example:  Molecular:  Complete ionic  Net ionic reaction

14. Day 2: Colligative Properties  The presence of solutes affects the properties of the solutions. Doesn’t matter which solute you have… just how many particles are present! These properties that depend on the [solute particles] are called colligative properties.

15. Colligative Properties 1.Vapor – pressure lowering 2.Boiling –point elevation 3.Freezing – point depression 4.Osmotic Pressure

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Colligative Properties of Electrolyte Solutions van’t Hoff factor, “i”, relates to the number of ions per formula unit. NaCl = 2, K 2 SO 4 = 3

17. van’t Hoff  van't Hoff was the first Nobel laureate in chemistry. He received the prize in 1901 to recognize his discovery of the laws of chemical kinetics, chemical equilibrium and osmotic pressure. He recognized the importance of concentration in determining reaction rates, emphasized the reversibility of reactions and treated equilibrium as a dynamic balance of opposing reactions. He invented the symbol still used to signify equilibria. van't Hoff showed the analogy between ideal gases and dilute solutions, and that osmotic pressure can be used to measure molecular weights in solution. Together with Wilhelm Ostwald, van't Hoff founded in 1887 the first journal devoted to physical chemistry, the Zeitschrift für Physikalische Chemie.  It is curious that the Nobel Prize was not given to him for his cornerstone proposal (independently reached by Jules Achille Le Bel) that the four bonds of carbon are directed to the corners of a regular tetrahedron, which explained optical isomerism. It was van't Hoff who, at age 22, published his revolutionary ideas that led chemists to think of molecules as objects with structures and 3-dimensional shapes. He constructed the first molecular models, of cardboard, to illustrate the principles of stereochemistry.  van't Hoff was born in Rotterdam, The Netherlands. Although he had strong interests in philosophy, poetry, mathematics and physics he studied organic chemistry with Kekulé; and Wurtz, then returned to the University of Utrecht for his Ph.D. (1874). In 1876 he became lecturer in physics at the Veterinary School in Utrecht and eventually (1878-96) Professor of chemistry, mineralogy, and geology at the University of Amsterdam. He disliked administrative work and teaching, and in 1896 accepted a professorship at the University of Berlin, with responsibility for just one lecture per week. His last publications applied mass-action principles to physiology and the investigation of enzyme reactions. Kekulé 

18. Practice van’t Hoff factor  Na 2 SO 4 i =  K 3 PO 4 i =  NaCl i =  (NH 4 ) 2 SO 4 i =  Mg(NO 3 ) 2 i =  C 12 H 22 O 11 i =

19. Figure 11.10: The presence of a nonvolatile solute inhibits the escape of solvent molecules from the liquid and so lowers the vapor pressure of the solvent.

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Boiling Point Elevation A nonvolatile solute elevates the boiling point of the solvent. A nonvolatile solute elevates the boiling point of the solvent.  T = i K b m solute K b = molal boiling point elevation constant m = molality of the solute

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Freezing Point Depression A nonvolatile solute depresses the freezing point of the solvent.  T = i K f m solute K f = molal freezing point depression constant m = molality of the solute

22. Copyright © Houghton Mifflin Company. All rights reserved. 11b–22 Figure 11.15: (a) Ice in equilibrium with liquid water. (b) Ice in equilibrium with liquid water containing a dissolved solute (shown in pink).

23. Copyright © Houghton Mifflin Company. All rights reserved. 11b–23 Figure 11.14: Phase diagrams for pure water (red lines) and for an aqueous solution containing a nonvolatile solute (blue lines).

24. Copyright © Houghton Mifflin Company. All rights reserved. 11b–24

25. The addition of antifreeze lowers the freezing point of water in a car’s radiator.

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Osmotic Pressure Osmosis: The flow of solvent into the solution through the semipermeable membrane. Osmotic Pressure: The excess hydrostatic pressure on the solution compared to the pure solvent.

27. Osmotic Pressure:    = i M R T  Where:   is the osmotic pressure  i is the van’t Hoff factor  M is molarity (moles/liter)  R is the gas constant (0.0821 L atm/K mole)  T is temperature in Kelvin

28. Really just Ideal Gas Law   = i M R T   = i moles R T Volume Volume  Volume = i moles RT P V = n RT

29. Figure 11.16: A tube with a bulb on the end that is covered by a semipermeable membrane.

30. Figure 11.17: The normal flow of solvent into the solution (osmosis) can be prevented by applying an external pressure to the solution. The minimum pressure required to stop the osmosis is equal to the osmotic pressure of the solution.

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 If the external pressure is larger than the osmotic pressure, reverse osmosis occurs. One application is desalination of seawater.

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Solution Composition 1.Molarity (M) = 2.Mass (weight) percent = 3.Mole fraction (  A ) = 4.Molality (m) =