1. Intermolecular Forces and Liquids and Solids Chapter 12

2. Midterm II Any conflicts with March 20? If yes, let me know ASAP. The original date was March 22.

3. Phase Diagram of Water Note the high critical temperature and critical pressure: –These are due to the strong van der Waals forces between water molecules. The slope of the solid–liquid line is negative. –This means that increasing the pressure above 1 atm will raise the boiling point (as expected) and lower the melting point. –Lower the melting point?

4. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures.

5. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures. At 1 atm, solid CO 2 does not melt at any temperature. Instead, it sublimes to form CO 2 vapor. Why might it be useful as a refrigerant?

6. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures. If you want to send something frozen across the country, you can pack it in dry ice. It will be frozen when it reaches its destination, and there will be no messy liquid left over like you would have with normal ice.

7. The slope of the curve between solid and liquid is positive for CO 2 as well as almost all other substances. Why does water differ?

8. Freeze-drying Completely remove water from some material, such as food, while leaving the basic structure and composition of the material intact Two reasons –Keeps food from spoiling for a long period of time –Significantly reduces the total weight of the food How? –Freeze the material –Lower the pressure (<0.006 atm) –Increase the temperature slightly Normal (right) and freeze-dried (left) spaghetti

9. Freeze-drying How? –Freeze the material –Lower the pressure –Increase the temperature slightly Normal (right) and freeze-dried (left) spaghetti

10. Physical Properties of Solutions Chapter 13

11. 13.1 A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount

12. A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature. An unsaturated solution contains less solute than the solvent has the capacity to dissolve at a specific temperature. A 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. 13.1

13. 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.

14. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them. If the solvent is water, then the process is called hydration.

15. How Does a Solution Form If an ionic salt is soluble in water, it is because the ion- dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

16. Energy Changes in Solution Simply, three processes affect the energetics of the process:  Separation of solute particles  Separation of solvent particles  New interactions between solute and solvent

17. Energy Changes in Solution The enthalpy change of the overall process depends on  H for each of these steps.

18. 13.2 Three types of interactions in the solution process: solvent-solvent interaction solute-solute interaction solvent-solute interaction  H soln =  H 1 +  H 2 +  H 3

19. “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) 13.2

20. Concentration Units The concentration of a solution is 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

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

22. 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 13.3 Strategy: Find mass of solvent Know mass of solute + mass of solvent = mass of solution If mass of solution and mass of solute known, can calculate mass of solvent Can calculate mass of solute from moles of solute Can calculate mass of solution from density and volume of the solution Solve

23. 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 5.86 moles of solute per 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

24. Temperature and Solubility Solid solubility and temperature solubility increases with increasing temperature solubility decreases with increasing temperature 13.4 No clear correlation between ΔH soln and the variation of solubility with temperature Stopped here