Homework for Unit 6. Liquids and aqueous solutions

Review of kinetic molecular theory of gases (Text 13.1)

1. List the four characteristics of an ideal gas.

2. a. What aspect of molecule behavior causes what we perceive as pressure?

b. What aspect of molecule behavior corresponds to what we perceive as temperature?

3. a. Compare the kinetic energies, masses, and velocities of CO2 molecules and H2 molecules at the same temperature.

b. How would changing the temperature affect the velocities of the gas molecules?

4. Explain each observation about gas behavior in terms of molecular motion. For example,

Observation: if you increase the volume of the container, the gas pressure decreases.

Explanation: the moving molecules fill the container; with more room the move around, the molecules collide with the container walls less often, so the pressure is lower.

a. If you heat the air in a sealed syringe, the gas expands (pushes the plunger out).

b. The air pressure inside a cold tire is less than inside a hot tire.

c. This reaction occurs in a rigid container at constant temperature:

2 NO (g) + O2 (g) → 2 NO2 (g)

The gas pressure is lower after the reaction than it was before the reaction occurred.

d. If you open a bottle of ammonia at one end of the room, after a while you can smell the ammonia at the other end of the room.

Sticking together

5. Under what conditions will a gas condense to form a liquid? Why do they do that?

6. How can we evaluate the strength of the attractive forces between molecules?

7. For molecules of similar structure, what generally happens to the boiling point as molar mass increases? What does this suggest about the attractive forces between molecules?

8. Predict which would probably have a higher boiling point, and why:

a. Ne or Ar b. Cl2 or O2

Phase change and temperature

9. a. Draw a heating curve for water. Label the states and phase changes. Indicate the melting point and boiling point on the graph (including numerical values).

b. Which phase changes are endothermic? Which are exothermic? State a general rule for predicting whether a phase change is endo- or exothermic.

10. a. Given that the heat of fusion of ice is about 6.0 kJ/mol, how much energy is needed to melt one ice cube that weighs 35 grams?

b. Given that the energy to warm liquid water is about 4.184 J/g °C, and the heat of vaporization of water is 41 kJ/mol, how much energy is needed to heat a cup of water (240 g water) from 25 °c to 100 °C, then convert all of the water to steam?

11. Why might a burn from steam at 100 °C be more serious than a burn from an equal amount of liquid water at 100 °C?

Phase change: liquid back to gas (pp 388-389 + section 13.2)

12. We have emphasized that at any given temperature, all molecules have the same average kinetic energy. Does that mean all molecules have the same kinetic energy?

13. Describe evaporation. Which molecules evaporate? Predict how evaporation would affect the temperature of the liquid.

14. Vapor pressure is the pressure caused by vapor in a closed container at equilibrium. What two processes are occurring at the same rate in this system? How is vapor pressure affected by temperature, and why?

15. What is the difference between evaporation and boiling? What determines the temperature at which a substance begins to boil? What is inside the bubbles in boiling water?

16. The graph at right shows the vapor pressures of five different substances at various temperatures.

a. At what temperature is the vapor pressure of substance (d) equal to 400 mm Hg?

b. What is the normal boiling point of substance (b)?

c. If the atmospheric pressure were 600 mm Hg, at what temperature would substance (c) boil?

d. Compare the average kinetic energies of substances (a) and (b) at room temperature (25 °C).

17. Why does it take longer to hard-boil an egg in Denver, the “mile-high city,” than it does in Fresno?

Solute + solvent = solution

18. A solution contains 2.70 g K2SO4 in 250.0 mL of solution. Identify the solute and solvent in this solution, then calculate its concentration in moles per liter.

19. How many moles of silver nitrate (AgNO3) are in 15.2 mL of 0.038 M AgNO3 solution?

20. How many mL of 0.0505 M HCl are needed to provide 0.001264 mol HCl?

21. What mass of NaHCO3 (baking soda) will neutralize 25 mL of 5.1 M H2SO4 (battery acid)?

H2SO4 (aq) + 2 NaHCO3 (s)  2 H2O (l) + 2 CO2 (g) + Na2SO4 (aq)

22. How many mL of 0.105 M KI solution will react with 25.0 mL of 0.0500 M Pb(NO3)2 solution?

2 KI (aq) + Pb(NO3)2 (aq)  PbI2 (s) + 2 KNO3 (aq)

23. In one experiment, 0.175 g CaCO3 reacts with 16.28 mL of HCl in this reaction:

2 HCl (aq) + CaCO3 (s)  CaCl2 (aq) + H2O (l) + CO2 (g)

Calculate the concentration of the HCl solution, in mol/L

24. How many L of H2 gas at STP form when 15.0 mL of 3.0 M HCl react?

2 Al (s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2 (g)

1