1. Chapter 11a : Solutions and Their Properties

2. Introduction 1. A mixture is any intimate combination of two or more pure substances 2. Can be classified as heterogeneous or homogeneous Heterogeneous -The mixing of components is visually nonuniform and have regions of different composition Homogenous -Mixing is uniform, same composition throughout -Can be classified according to the size of their particles as either solutions or colloids

3. What is a solution? Sugar + water

4. Solution Formation Solute Dissolved substance, or smaller quantity substance Solvent Liquid dissolved in, larger quantity substance Saturated solution Contains the maximum amount of solute that will dissolve in a given solvent.

5. Solution Formation Unsaturated Contains less solute than a solvent has the capacity to dissolve. Supersaturated Contains more solute than would be present in a saturated solution. Crystallization The process in which dissolved solute comes out of the solution and forms crystals.

6. Types of Solutions

7. Solutions Solution 1. Homogeneous mixtures 2. Contain particles with diameters in the range of 0.1–2 nm 3. Transparent but may be colored 4. Do not separate on standing Colloids 1. Milk & fog 2. Diameters 2-500 nm 3. Do not separate on standing

8. Types of colloids Aerosol – liquid in gas Solid Aerosol – solid in gas

9. Emulsion – liquid in liquid like oil droplets in mayonnaise. Foams – gases in liquids like whipped cream Solid emulsion – liquid in a solid like milk in butter

10. Energy Changes and the Solution Process Three Types of interactions 1. Solvent-solvent 2. Solvent-solute 3. Solute-solute “Like dissolves like” solutions will form when three types of interactions are similar in kind and magnitude

11. Energy Changes and the Solution Process Example NaCl and water: Ionic solid NaCl dissolve in polar solvents like water because the strong ion-dipole attractions between Na + and Cl - ions and polar water molecules are similar in magnitude to the strong dipole-dipole attractions between water molecules and to the strong ion-ion attractions between Na + and Cl - ions Example Oil and water Oil does not dissolve in water because the two liquids have different kinds of intermolecular forces. Oil is not polar or an ionic solvent

12. Energy Changes and the Solution Process NaCl in H 2 O 1. Ions that are less tightly held because of their position at a corner or an edge of the crystal are exposed to water molecules 2. Water molecules will collide with the NaCl until an ion breaks free 3. More water molecules then cluster around the ion, stabilizing it by ion- dipole attractions 4. The water molecules attack the weak part of the crystal until it is dissolved 5. Ions in solution are said to be solvated they are surrounded and stabilized by an ordered shell of solvent molecules

13. Energy Changes and the Solution Process  G, Free energy change 1. If  G is negative the process is spontaneous, and the substance is dissolved 2. If  G is positive the process is non-spontaneous, the substance is not dissolved 3.  G =  H -T  S  H, enthalpy, heat flow in or out of the system,  H soln heat of solution  S, entropy, disorder,  S soln entropy of solution

14. Energy Changes and the Solution Process

15.  S soln Entropy of Solution Usually a positive number because when you dissolve something you are increasing disorder  H soln Heat of Solution 1. Harder to predict because it could be exothermic (-  H soln ) or endothermic (+  H soln ) 2. The value of the heat of solution for a substance results from an interplay of the three kinds of interactions

16. Energy Changes and the Solution Process 1)Solvent-solvent interactions: Energy is required (endothermic) to overcome intermolecular forces between solvent molecules because the molecules must be separated and pushed apart to make room for solute particles 2)Solute-solute interactions: Energy is required (endothermic) to overcome interactions holding solute particles together in a crystal. For an ionic solid, this is the lattice energy. Substances with higher lattice energies therefore tend to be less soluble than substances with lower lattice energies.

17. Energy Changes and the Solution Process 3) Solvent-solute interactions: Energy is released (exothermic) when solvent molecules cluster around solute particles and solvate them. For ionic substances in water, the amount of hydration energy released is generally greater for smaller cations than for larger ones because water molecules can approach the positive nuclei of smaller ions more closely and thus bind more tightly. Hydration energy generally increases as the charge on the ion increases.

18. Energy Changes and the Solution Process The solute–solvent interactions are stronger than solute–solute or solvent–solvent. Favorable process Exothermic rxn. Exothermic -  H soln

19. Energy Changes and the Solution Process The solute–solvent interactions are weaker than solute–solute or solvent–solvent. Unfavorable process. Endothermic rxn Endothermic +  H soln

20. Energy Changes and the Solution Process Hydration –The attraction of ions for water molecules Hydration Energy –The energy associated with the attraction between ions and water molecules Lattice energy –The energy holding ions together in a crystal lattice

21. Example 1 Arrange the following in order of their expected increasing solubility in water: Br 2, KBr, C 7 H 8 C 7 H 8 is non-polar and is insoluble in water Br 2 is non-polar but because of its size its polarizable and is soluble KBr is an ionic compound and is very soluble in water

22. Example 2 1.Na+ 2.Cs+ 3.Li+ 4. Rb+ 1.Mg 2+ 2.Na + 3.Li + Which would you expect to have the larger (more negative) hydration energy?

23. Summary Solution is a homogeneous mixture, which consists solute and solvent. Like dissolves like Energy changes in solution process: A)  S>0 B)  H can be either endothermic or exothermic