1. Metallic Solids Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces In metals valence electrons are delocalized throughout the solid “Electron-sea model”

2. Types of Bonding in Crystalline Solids

3. Vapor Pressure and Boiling Point Boiling point ≡ temperature at which vapor pressure equals atmospheric pressure. Normal boiling point ≡ temperature at which vapor pressure is 760 torr. Fig 11.24

4. Physical Properties of Solutions Chapter 13

5. Solution - a homogenous mixture of 2 or more substances Solute - the substance(s) present in the smaller amount(s) Solvent - the substance present in the larger amount Table 13.1

6. Solutions The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

7. © 2009, Prentice- Hall, Inc. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them. The Effect of Intermolecular forces Fig 13.1 Dissolution of an ionic solid

8. Fig 13.2 Hydrated Na + and Cl − ions The negative end of the water dipoles point toward the positive ion Positive ends point toward the negative ion Result is hydrated ions

9. Three types of interactions in the solution process: solute-solute interaction solvent-solvent interaction solvent-solute interaction  H soln =  H 1 +  H 2 +  H 3 Fig 13.2

10. Energy Changes in Solution Formation The enthalpy change of the overall process depends on  H for each of these steps. Fig 13.4 Enthalpy changes accompanying solution processes

11. Solutions Solution Formation, Spontaneity, and Entropy Enthalpy is only part of the picture Increasing the disorder or randomness of a system tends to lower the energy of the system Entropy ≡ degree of randomness or disorder in a system Solutions favored by increase in entropy that accompanies mixing Fig 13.6

12. Caveat Emptor! Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved. Dissolution is a physical change — you can get back the original solute by evaporating the solvent If you can’t, the substance didn’t dissolve, it reacted. Fig 13.7

13. Unsaturated solution - contains less solute than the solvent has the capacity to dissolve at a specific temperature Supersaturated solution - contains more solute than is present in a saturated solution at a specific temperature Sodium acetate crystals rapidly form when a seed crystal is added to a supersaturated solution of sodium acetate. Saturated solution - contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature Fig 13.10

14. “like dissolves like” Two substances with similar intermolecular forces are likely to be soluble in each other: non-polar molecules are soluble in non-polar solvents CCl 4 in C 6 H 6 polar molecules are soluble in polar solvents C 2 H 5 OH in H 2 O ionic compounds are more soluble in polar solvents NaCl in H 2 O or NH 3 (l) Factors Affecting Solubility

15. Acetone is miscible in water H2OH2O C 6 H 14

16. Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water Cyclohexane (which only has dispersion forces) is not Fig 13.12 Structure and solubility

19. Pressure Effect on Gases in Solution Solubility of liquids and solids does not change appreciably with pressure Solubility of a gas in a liquid is directly proportional to its pressure Fig 13.14 Effect of pressure on gas solubility

20. Henry’s Law S g = kP g where S g ≡ solubility of the gas k ≡ the Henry’s Law constant for that gas in that solvent P g ≡ partial pressure of the gas above the liquid Fig 13.15 Solubility decreases as pressure decreases

21. Temperature Effect on Solids and Liquids Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature Fig 13.17 Solubilities of several ion compounds as a function of temperature

22. The opposite is true of gases: Carbonated soft drinks are more “bubbly” if stored in the refrigerator Warm lakes have less O 2 dissolved in them than cool lakes Temperature Effect on Gases Fig 13.18 Variation of gas solubility with temperature

24. Concentration Units Concentration - the amount of solute present in a given quantity of solvent or solution. Percent by Mass % by mass = x 100% mass of solute mass of solute + mass of solvent = x 100% mass of solute mass of solution 13.3 Mole Fraction (X) X A = moles of A sum of moles of all components

25. Concentration Units Continued M = moles of solute liters of solution Molarity (M) Molality (m) m = moles of solute mass of solvent (kg) 13.3

26. What is the molality of a 5.86 M ethanol (C 2 H 5 OH) solution whose density is 0.927 g/mL? m =m = moles of solute mass of solvent (kg) M = moles of solute liters of solution Assume 1 L of solution: 5.86 moles ethanol = 270 g ethanol 927 g of solution (1000 mL x 0.927 g/mL) mass of solvent = mass of solution – mass of solute = 927 g – 270 g = 657 g = 0.657 kg m =m = moles of solute mass of solvent (kg) = 5.86 moles C 2 H 5 OH 0.657 kg solvent = 8.92 m 13.3