Solutions

1. Solutions Chapter 14

2. Key concepts • Understand the solvation process at the molecular level. • Be able to qualitatively describe energy changes during solution formation. • Understand how entropy affects the solution process. • Know the terms solute, solvent, saturated, unsaturated, and supersaturated. • Know factors contributing to miscibility and immiscibility. • Know how pressure affects solubility of gases: Henry’s law • Understand the effects of colligative properties: Raoult’s Law, freezing/boiling point changes, and osmosis. • Know what a colloidal suspension is and how it may differ from a solution.

3. Solution process • When intermolecular attractions between solute molecules are overcome by intermolecular attractions with solvent. • solute breaks up to become dissolved within solvent. • process of solvation

4. Solids and liquids Overcoming intermolecular forces between solvent molecules (h-bonds in water)

5. Solids and liquids Overcoming intermolecular forces between solute molecules/ions

6. Solids and liquids Solvent molecules surround and interact with solute (solvation) Solvation energy (hydration energy) = step b + step c

7. Solvation and heat • The heat flow associated with formation of a solution is called the heat of solution (Hsolution) • If solvation process is exothermic, then Hsolution is ___________. • If solvation process is endothermic, then Hsolution is ___________. • DEMO: Ammonium nitrate vs calcium chloride

8. Two thermodynamic principles of solution formation • spontaneous processes tend to be exothermic (Hsolution is ?) • processes where entropy increases tend to occur spontaneously • solution formation favored by an increase in entropy • Solution formation may involve an endothermic process IF the entropy change outweighs the need to absorb heat.

9. Liquid-liquid interactions • solute-solvent interactions: like dissolves like • miscible: • immiscible: • Molecular “red rover”…

10. Dissolving gases in liquids • Miscibility rules for liquids also apply to gas/liquid solutions. • Usually, aqueous solutions of gases occur when • Hydrogen bonding is possible with the gas • The gas ionizes extensively in water • Rxn between gas and water takes place

11. dynamic equilibrium of solutions • solvent + solute  solution • saturated solution: in equilibrium with undissolved solute • unsaturated solution: not enough solute available to reach equilibrium.

12. crystallization • for crystallization to take place, two conditions must be met: • solution must reach saturation point • crystal order must be established • Supersaturated solution: an unstable solution with a concentration higher than the equilibrium concentration. • Solution is beyond normal saturation point, but a crystal order has not yet established.

13. Temperature effects on solubility • Increasing temperature may increase OR decrease solubility. Generally… • Solubility of substances that undergo endothermic dissolution ________ with increasing temperature • Solubility of substances that undergo exothermic dissolution ________ with increasing temperature

14. Dissolving gases in solution: why your pop bottle fizzes over…. • The solubility of a gas is directly proportional to partial pressure of the gas. • if pressure on gas increases, the solubility of the gas also increases. • Henry’s law • Pg = k Cg • solubility and temperature:

15. Colligative properties • Physical properties of solutions that depend on the total number of solutes (not type of solute) are colligative properties. • We will look at several types of colligative properties.

16. Raoult’s Law: solutes and vapor pressure • Partial pressure of a liquid solvent is dependent upon the amount of solute it contains. Psolvent = Xsolvent P0solvent • This form of the law applies only to non-volatile, non-ionizing solutes in ideal solutions.

17. Raoult’s law for two volatile liquids • Two volatile liquids mutually effect each other’s partial pressures. • PA = XAP0A • PB = XBP0B

18. Deviations from ideality • Positive and negative • Depend on interactions between the two substances. • Mixtures of polar/non-polar substances usually lead to _______ deviations of the ideal pressure. • Mixtures of polar substances usually lead to ________ deviations of the ideal pressure.

19. Fractional distillation • A process by which two volatile liquids are separated. • Each time boiling and condensation takes place, the resulting liquid contains an increasing amount of the more volatile component. • Fractional distillation columns are designed to facilitate this process.

20. Boiling point elevation/freezing point depression • T = kf m or • T = kb m • T is change in temperature • Must include molality of all solutes in solution. • Kf/b constants have been determined for several substances… (p 556) • What is the change in freezing point temperature for a 0.100 molal solution of sucrose in water? • What is the change in freezing point temperature for a 0.100 molal solution of K2CO3 in water?

21. Ion association • Ions in solution occasionally “stick” to each other. When this happens, the combination behaves as one particle in solution. • The “effective concentration” of ions in solution depends on the amount of association that occurs.

22. Physical behavior of ionic solutions Non-ideal ionic solutions Near ideal ionic solutions H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O H2O

23. van’t Hoff factor • Table 14-3 (p. 561). What happens as the number of ions per formula unit increases? • What will the actual Tf be for a 0.100 m solution of K2CO3?

24. Osmotic pressure: making pickles h Height of column  pressure head from osmosis

25. Calculating osmotic pressure • For dilute solutions, the osmotic pressure is given by • = MRT • Example: 0.100 molar solution of sucrose in water at 20 C. • Osmotic pressure plays a key role in biological processes.

26. Molecular mass of polymers • A 2.30 g sample of a polymer is dissolved in 250 mL of water, resulting in an osmotic pressure of 5.40 torr. • What is the molecular mass of the polymer?

27. Reverse osmosis • If a pressure is applied to the solution side of the system, greater than the osmotic pressure (2.40 atm in our sugar example), pure liquid can be forced out of the membrane. • Everything moves in the opposite direction • Desalination of seawater—26 atm pressure required.

28. Colloids • Colloids, or colloidal suspensions, are in an “in-between” area between homogeneous and heterogenous mixtures. • Solute-like particles are dispersed in a solvent-like medium. • Colloidal suspensions disperse and scatter light (Tyndall effect) • Example: milk.

29. Hydrophilic colloids • Hydrophilic colloids form between water and dispersed phases containing polar groups about their surfaces. • Example: hemoglobin

30. Hydrophobic colloids • Hydrophobic dispersed particles are not attracted to water. In order to produce a colloidal suspension, we must introduce an emulsifying agent. • How to make your own mayo…. • Soaps are excellent examples of emulsifying agents.

31. Micelles • Contain a “polar head” and a “non-polar tail”. • Non-polar side attracts to non-polar substances (dirt, grease, etc). • Polar side attracts to water. • Micelles surround non-polar particle, suspending it in aqueous solution.