CIS Chemistry II Review Chapters 11 & 13

CIS Chemistry II Review – Chapters 11 & 13

1. What type(s) of intermolecular force is (are) common to

a) Xe and methanol (CH3OH) b) CH3OH and CH3CN c) NH3 and HF

2. Describe the intermolecular forces that must be overcome to convert each of the following from a liquid

to a gas: a) N2 b) C2H5OH c)H2Se

3. Account for the steady increase in boiling point of the noble-gas elements with increasing atomic weight (Table 11.3).

4. Rationalize the difference in boiling points between the members of the following pairs of substances: a) HF (20ºC) and HCl ( -85ºC)

b) Br2 (59ºC) and ICl ( 97ºC)

5. Out of the following substances, which would you expect to be the most volatile?

CH3OH CH4 C2H6

6. For a given substance, how does ΔHfus usually compare to the ΔHvap? Why?

7. Using the definition of boiling, describe the effect pressure has on the boiling point of a liquid. 8. Consider the phase diagrams for water and carbon dioxide:

a. Describe all the phase changes that would occur if water originally at 100.0ºC and 0.50 atm is cooled at constant pressure until the temperature is -10ºC.

b. Describe all the phase changes and the temperatures at which they occur when carbon dioxide is heated from -80ºC to -20ºC at a constant pressure of 6 atm. c. In the phase diagram involving water, the slope of the solid-liquid equilibrium curve has a negative slope, whereas the same line on the carbon dioxide phase diagram has a positive slope. Explain the significance of this observation.

9. Explain the following observations: a) Water evaporates more quickly on a hot, dry day than on a hot, humid day. b) It takes longer to cook hard-boiled eggs at high altitudes than at lower altitudes.

10. What is the difference between a crystalline and an amorphous solid?

11. Consider the following substances in their solid states and classify each according to the type of solid:

Au Quartz FeCl3 C6H12O6 graphite

12. What are the two major components of any solution?

13. The enthalpy of solution formation is the sum of three steps. Describe what happens in each step, and state whether each step is endo or exothermic. 14. Explain the statement “like dissolves like.”

15. Why is ethanol more soluble in water than pentanol?

16. How does the solubility of a gas change as we increase the partial pressure of that gas above the liquid?

17. Calculate the concentration of CO2 in a soft drink after the bottle is opened and equilbriates at 25ºC under -4 a CO2 partial pressure of 3.0 x 10 atm. The Henry’s law constant for CO2 in water at this temperature is 3.1 x 10-2 mol/L-atm.

18. State how the temperature affects the solubility of most a) solids; b) gases.

19. Calculate the mole fraction and molality of ethanol in a solution that contains 100.0 grams of water and

100.0 grams of ethanol (C2H5OH).

20. The molal freezing point depression constant, Kf, is 1.86 ºC/m. A solution of 22.1 grams of a nonelectrolyte was dissolved in 750.0 grams of water and the freezing point of the solution was observed to be -2.35ºC. Determine the molecular weight of the substance.

21. From the following list of substances, KBr (s); C2H5OH (l); CO2 (g); C6H12 (l); Select the one that best satisfies each of the following descriptions. Give a brief explanation for each choice. A. Likely to mix in all proportions with water B. Solubility in water is strongly affected by pressure C. Produces a conducting solution when dissolved in water D. Has the lowest solubility in water 22. A solution that is prepared by dissolving 3.15 grams of an unknown substance in 25.0 grams of benzene,

C6H6, has a freezing point of 1.12ºC. The normal freezing point of benzene is 5.50ºC, and Kf for benzene is 5.12ºC•kg/mol. A. Calculate the molar mass of the unknown substance. B. Calculate the mole fraction of benzene in the solution described above. C. If the vapor pressure of pure benzene at 35ºC is 150.0 mm Hg, calculate the vapor pressure of the solution described above at 35ºC.

23. An aqueous solution contains 4.50 grams of NaCl in 750.0 mL of solution. Calculate the osmotic pressure of the solution at 22.0ºC.

24. List the following aqueous solutions in order of their expected freezing points: 0.050 m MgCl2; 0.15 m KBr; 0.10 m HBr; 0.050 m HC2H3O2; 0.10 m sugar

25. Seawater contains 0.0079 g Sr2+ per kilogram of water. What is the concentration of Sr2+ measured in ppm?